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Chapter 13 Influence of Phosphate on Adsorption and Surface Precipitation of Lead on Iron Oxide Surfaces
Developments in Earth & Environmental Sciences, 7 Mark O. Barnett and Douglas B. Kent (Editors) r 2008 Elsevier B.V. All rights reserved DOI 10.1016/S1571-9197(07)07013-9
Chapter 13
Influence of Phosphate on Adsorption and Surface Precipitation of Lead on Iron Oxide Surfaces Liyun Xie and Daniel E. Giammar Department of Energy, Environmental & Chemical Engineering, Washington University in St. Louis, One Brookings Dr. St. Louis, MO 63130, USA
ABSTRACT Lead and phosphate sorption on goethite-coated and uncoated quartz sand was measured experimentally and modeled within a reaction-based framework. Single sorbate batch experiments and experiments with lead and phosphate present together were conducted. Adsorption of lead and phosphate to goethite-coated sand was dominated by adsorption to the goethite coating. A surface complexation model adapted from models for pure goethite successfully simulated lead and phosphate adsorption to goethite-coated sand over a broad range of pH and total sorbate concentrations; the inclusion of a surface complexation reaction for lead adsorption to the quartz surface was necessary to improve the model fit. Lead sorption on goethite-coated sand was enhanced by the presence of phosphate. The effect was most pronounced at low pH. The enhanced lead uptake was predicted by the combination of the single sorbate surface complexation models. The adsorption of phosphate at low pH decreased the surface charge and potential, which increases the extent of lead adsorption. No ternary surface complexes were needed to model the dual sorbate results. When a reaction for the precipitation of chloropyromorphite (Pb5(PO4)3Cl) was included, the model predicted precipitation only at the lowest pH and highest phosphate loading studied. Over most experimental conditions, including conditions that were initially supersaturated with respect to chloropyromorphite, the equilibrium model predicted that adsorption was the dominant mechanism of lead sorption; however, the actual mechanisms may be controlled by the relative rates of precipitation and adsorption reactions.
Corresponding author. Tel.: +314-935-6849; Fax: +314-935-5464;
E-mail: [email protected] (D.E. Giammar).
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13.1. Introduction The fate and transport of lead in soil and groundwater is strongly controlled by chemical reactions at the mineral–water interface. Surface reactions include adsorption and precipitation. Goethite (a-FeOOH), a common iron(III) oxyhydroxide mineral in natural systems, and other iron(III) oxides and oxyhydroxides are important environmental sorbents for heavy metals due to their high specific surface areas and reactive surfaces (Dzombak and Morel, 1990; Coston et al., 1995; Schwertmann and Cornell, 2000). Iron oxide coated sands prepared in the laboratory have previously been used as model materials for real soils and sediments that have iron oxide coatings (Gabriel et al., 1998; Cheng et al., 2004). Lead adsorbs to goethite through the formation of inner-sphere surface complexes with surface hydroxyl groups (RFeOH) as determined by extended X-ray adsorption fine-structure spectroscopy (EXAFS) (Roe et al., 1991; Bargar et al., 1997). Lead also adsorbs via inner-sphere complexation as a mixture of monodentate and bidentate complexes on ferrihydrite (Fe5HO8 4H2O) (Trivedi et al., 2003) and as mononuclear bidentate complexes on hematite (a-Fe2O3) (Bargar et al., 1997). Evidence for inner-sphere complexation is also provided by a lack of ionic strength effects on adsorption. The ionic strength of the bulk solution has less impact on inner-sphere complexation than on outer-sphere complexation (Hayes and Leckie, 1986). Adsorption of lead to goethite has been successfully interpreted through the inclusion of reaction 1 in a surface complexation model (SCM) (Hayes and Leckie, 1986). FeOH þ Pb2þ ¼ FeOPbþ þ Hþ
(13.1)
Other SCMs have included reaction 1 as well as two other reactions (2 and 3) (Gunneriusson et al., 1994). FeOH þ Pb2þ ¼ FeOHPb2þ
(13.2)
FeOH þ Pb2þ þ H2 O ¼ FeOPbOH þ 2Hþ
(13.3)
Surface complexation modeling considers the adsorption of inorganic species to specific reactive surface sites and includes terms to account for both the chemical and electrostatic energetics of adsorption (Dzombak and Morel, 1990). The model of Hayes and Leckie (1986) successfully predicted the pH dependence of lead adsorption. Although metal adsorption to quartz is not usually significant in soil and groundwater relative to adsorption to iron oxyhydroxides and clays, metal
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cations can still adsorb at the quartz surface. X-ray absorption spectroscopy investigations found inner-sphere lead surface complexes formed on quartz (Chen et al., 2006) and amorphous silica (Elzinga and Sparks, 2002). Lead sorption at the iron oxide–water interface can be enhanced by the presence of phosphate. Enhancement can occur as the result of: (a) changes to surface charge that make lead adsorption more favorable, (b) formation of ternary lead–phosphate–iron oxide surface complexes, (c) precipitation of lead phosphate solids, and (d) surface alteration from the formation of an iron(III) phosphate surface precipitate. Phosphate addition to leadcontaminated soils has been proposed as a means of reducing lead mobility and bioavailability (Ruby et al., 1994; Ma et al., 1995; Hettiarachchi et al., 2000; Yang et al., 2001; Cao et al., 2002; Ryan et al., 2004). Although remediation strategies usually suggest using phosphate addition to precipitate lead phosphate solids, enhancement of lead binding to iron oxyhydroxides may also increase lead sorption even without lead phosphate precipitation. Phosphate adsorbs to goethite through inner-sphere complex formation as observed directly with attenuated total reflectance Fourier transform infrared spectroscopy (Arai and Sparks, 2001). Atomic force microscopy (AFM) of phosphate adsorption to the (0 1 0) surface of a goethite single crystal observed phosphate adsorbed in a 1:1 ratio with the singly coordinated hydroxyl groups via monodentate complexes (Dideriksen and Stipp, 2003). A combination of three monodentate surface complexes (reactions 4–6) has been proposed for modeling phosphate adsorption to goethite (Nilsson et al., 1992). þ FeOH þ H2 PO 4 þ H ¼ FePO4 H2 þ H2 O
(13.4)
FeOH þ H2 PO 4 ¼ FePO4 H þ H2 O
(13.5)
2 þ FeOH þ H2 PO 4 ¼ FePO4 þ H þ H2 O
(13.6)
When metal ions (e.g., M2+) and anions (e.g., L) adsorb together on the oxide surface, ternary surface complexes can form by either metal-bridging (reaction 7) or ligand-bridging (reaction 8) reactions (Hering and Kraemer, 1994). SOH þ M2þ þ L ¼ SOM L þ Hþ
(13.7)
SOH þ L þ M2þ þ Hþ ¼ SLM2þ þ H2 O
(13.8)
Ternary surface complexation is an adsorption process with the simultaneous accumulation of metal cations and ligand anions at the solid– water interface without the development of three-dimensional structure
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(Sposito, 1986). This adsorption process is in contrast to the formation of a precipitate at the surface that has longer range three-dimensional structure. Lead and phosphate can also precipitate at the goethite surface, most likely through heterogeneous nucleation of a lead phosphate precipitate on the goethite substrate. Once precipitation is initiated, the number of surface sites for lead or phosphate binding is no longer fixed by the goethite surface area. Surface precipitation is most likely at high sorbate loading. Potential lead-containing solids include lead hydroxide (Pb(OH)2) and the lead phosphates chloropyromorphite (Pb5(PO4)3Cl), hydroxypyromorphite (Pb5(PO4)3OH), and lead hydrogen phosphate (PbHPO4). Ler and Stanforth (2003) observed precipitation of lead and phosphate on goethite in a phase with the stoichiometry of chloropyromorphite or hydroxypyromorphite. The reaction of hydroxyapatite (Ca5(PO4)3OH) and goethite with adsorbed lead resulted in the formation of chloropyromorphite through homogeneous nucleation in the solution and possibly heterogeneous nucleation at the goethite surface (Zhang et al., 1997). In this study we examine the effects of phosphate on lead sorption to goethite-coated sand. The word sorption is used here to describe all processes (adsorption, surface precipitation, co-precipitation) that result in accumulation of lead at the goethite–water interface. The objectives of this project were to determine the utility of goethite-coated sand as a model porous medium with a reactive surface phase, distinguish the different mechanisms through which phosphate can increase lead sorption, and determine the geochemical conditions under which each mechanism is dominant. Once the interfacial chemical reactions on goethite-coated sand have been established, it can be a useful material for evaluating the effects of surface reactions on lead transport. The different mechanisms of lead sorption in the presence of phosphate may result in similar equilibrium partitioning of lead between the solid and dissolved phases, but the rates of transport are likely to be affected by differences in sorption mechanisms.
13.2. Experimental 13.2.1. Materials All chemicals used were certified ACS grade (Fisher Scientific) unless otherwise specified. The concentrated nitric acid was trace metal grade. Pure goethite was synthesized from an alkaline aqueous system according to methods of Schwertmann and Cornell (2000). A 100 ml volume of 1 M
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Fe(NO3)3 was mixed with 180 ml of 5 M KOH in a 2 l polyethylene bottle and then immediately diluted with ultrapure water to 2 l. The bottle was sealed and heated in an oven at 651C for 60 h. An ocher precipitate formed during the heating process. The resulting suspension was washed free of excess dissolved ions, primarily K+ and NO 3 , by dialysis. The dialyzed suspension was freeze dried. The identity of the synthetic goethite (a-FeOOH) was confirmed by X-ray powder diffraction (XRD) (Rigaku Geigerflex D-MAX/A Diffractometer) with Cu–Ka radiation. Solid samples for XRD were ground to powders and loaded into glass sample holders for analysis. Quartz sand (U.S. Silica) with particle sizes of 212–355 mm was soaked in 1 M HCl (Fisher Scientific) overnight to remove surface impurities from the quartz. The sand was rinsed with ultrapure water to remove the acid and then dried at 1051C. Goethite-coated sand was then prepared by agitating a mixture of acid-washed quartz sand and goethite at pH 6.8 and 0.01 M ionic strength (NaNO3) for 24 h. These conditions had previously been determined to be optimal for goethite attachment based on the surface charges of goethite and quartz (Scheidegger et al., 1993). Unattached goethite was removed by repeatedly suspending the sand in 0.01 M NaNO3 solution, allowing the sand to settle, and decanting the supernatant. These steps were repeated until a clear supernatant was achieved. The iron content of the coated sand was determined by extraction in a solution of 0.3 M sodium dithionite, 0.3 M sodium citrate, and 0.2 M sodium bicarbonate. This solution effectively dissolves crystalline Fe(III) oxides and oxyhydroxides by reducing them to soluble Fe(II) species (Clark et al., 1996). The extraction mixture was shaken for 1 h, set overnight at room temperature for complete reaction, and the extract was then analyzed by inductively coupled plasmaoptical emission spectrometry (ICP-OES). The goethite-coated quartz sand contained 0.2% (w/w) iron; the value of 0.005% iron in uncoated sand was reported by the supplier (U.S. Silica). The typical crystalline morphology of goethite and of goethite present on goethite-coated sand was observed with a field emission scanning electron microscope (Hitachi S-4500) (Fig. 13.1). Dry samples were fixed to sample holders with double-sided tape. To improve the SEM images, gold coating was used on some samples to increase electrical conductivity. The quartz sand surface was partially covered by goethite particles but large portions of quartz remained exposed. Table 13.1 summarizes the properties of pure goethite, uncoated quartz sand, and goethite-coated sand. The size of uncoated and goethite-coated sand were the same, but the specific surface area of the coated sand was more than three times that of uncoated sand due to the high surface area goethite coating. Specific surface area (m2/g) was measured by
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Figure 13.1: Scanning Electron Micrographs of Synthetic Goethite (Left) and Goethite-Coated Sand (Right).
Table 13.1: Sorbent Properties.
Size (mm) Iron (wt%) Specific surface area (m2/g)
Goethite
Quartz sand
Goethite-coated sand
1 62.8 33.6
212–355 0.005 0.05
212–355 0.2 0.16
BET-N2 adsorption (Quantachrome Autosorb AS-1). Prior to BET-N2 measurement, wet samples were freeze dried (Labconco, Freezone 4.5) and then loaded into sample cells for analysis. Based on the iron content of the goethite-coated sand and assuming the additional 0.11 m2/g of surface area for the coated sand is from goethite, the goethite in the coating is calculated to have a specific surface area of 35 m2/g, which is very close to the value of 33.6 m2/g measured for the pure goethite.
13.2.2. Recirculating Micro-Column Batch Adsorption Experiments Lead and phosphate sorption reactions were studied by continuously recirculating solution from a 50 ml reservoir through 1.8 ml columns that were filled with 2.5 g of either goethite-coated sand or uncoated sand. Columns with recirculation, instead of stirred batch reactors, were used because goethite detachment from quartz sand was observed when suspensions of goethite-coated sand were agitated on a rotary shaker or with a magnetic stir bar. The recirculation flow rate was controlled at 0.5 ml/min with a peristaltic pump. Samples were collected after 24 h, which allowed sufficient
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residence times of both the column (3 min) and the reservoir (100 min) to allow the entire system to be considered as one well-mixed batch reactor. Lead adsorption to goethite occurs on time scales much faster than those of this experiment (Hayes and Leckie, 1986), and our preliminary experiments demonstrated that 24 h was also sufficient for equilibration. After the 24 h equilibration period, a 20 ml effluent sample was collected for pH and flow-rate measurement and another 10 ml sample was collected for dissolved lead and/or phosphate analysis. For many of the single sorbate experiments the influent reservoir was then refilled with 30 ml of solution to provide a new desired composition and the system was equilibrated for another 24 h until the next sampling time. By using this approach, multiple adsorption datapoints were generated while using a limited amount of solid material and maintaining a completely saturated environment in the columns. Single sorbate (lead-only or phosphate-only) and binary sorbate (lead and phosphate together) experiments were conducted (Table 13.2). Single sorbate experiments were conducted to evaluate the extent to which adsorption to goethite-coated sand could be explained by adsorption to pure goethite. The influent lead and phosphate concentrations were undersaturated with respect to lead hydroxide and iron phosphate precipitates. The range of total lead concentrations was 107–104 M and that of total phosphate concentrations was 107–5 105 M. Lead was added as Pb(NO3)2 (Acros Organic) and phosphate as KH2PO4 (Fisher Scientific). The desired pH of each adsorption experiment was from 4 to 8 and was controlled by MES (Sigma) or HEPES (Acros Organic) buffer solutions adjusted to pH 6–8 by addition of NaOH. These buffers were selected because of their established low affinities for metal complexation (Good et al., 1966; Soares et al., 1999). Buffer concentrations were 1 mM for lead adsorption and 10 mM for phosphate adsorption experiments. The buffer concentrations were selected to be at least an order of magnitude higher than the maximum total adsorbate concentration studied, and phosphate was studied at a higher total concentration than was lead. For control of the ionic strength, solutions also contained 1 mM NaNO3 for lead adsorption experiments and 10 mM with NaNO3 for phosphate adsorption experiments. When combined with the buffer species and associated NaOH was added, the total ionic strength of the solutions was 1.2–1.8 mM for lead experiments and 12–18 mM for phosphate experiments, depending on the pH. To examine the transition between adsorption and surface precipitation, binary sorbate experiments were conducted in which lead and phosphate were both present. The pH in these experiments was kept at 4–7 and the initial lead concentration was fixed at 5 107 M. To provide the desired
356
Experiment no.
A-1 A-2 B-1 B-2 C-1 C-2 a
Sorbent
Uncoated sand Goethite-coated sand Uncoated sand Goethite-coated sand Uncoated sand Goethite-coated sand
Solution composition pH
[Pb]init (M)
[P]init (M)
[Cl]init (M)
SICPY,
4–8 4–8
107–104a 107–104a
0 0
0 0
NA NA
4–8 4–8
0 0
107–5 105a 107–5 105a
0 0
NA NA
4–7 4–7
5 107 5 107
7.5 1010–103 7.5 1010–103
2.5 1010–104 2.5 1010–104
3, 1, 1, 5 3, 1, 1, 5
Single columns were used to evaluate adsorption over a broad concentration range at a given pH for lead-only and phosphate-only experiments.
init
L. Xie and D. E. Giammar
Table 13.2: Conditions of Experiments Conducted.
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initial saturation index with respect to chloropyromorphite, phosphate concentrations were varied from 7.5 1010 to 1.2 103 M and chloride concentrations from 2.5 1010 to 4 104 M (Table 13.3). These values were selected to approach relevant environmental concentrations while avoiding the use of initial solutions that were at even higher initial conditions of saturation. The initial saturation index (SI) (Eq. (13.9)) was examined at two conditions of undersaturation (3 and 1) and two of supersaturation (+1 and +5). The solubility product (Ksp) of chloropryromorphite is 1084.43 (Nriagu, 1973). 3 fPb2þ g5 fPO3 4 g fCl g SI ¼ log (13.9) K sp At negative values of SI, the precipitation of chloropyromorphite is not expected, and at positive SI values, chloropyromorphite precipitation is thermodynamically favorable. In sorbent-free control experiments, supersaturated conditions could be maintained for times of more than a day without loss of lead or phosphate from solution; however, at longer time scales, chloropyromorphite precipitation was observed for supersaturated conditions. Only the initial saturation indices were set at fixed values, and once lead and phosphate are taken up by the solid, either by adsorption or precipitation, the saturation index will decrease from its initial value. All experiments were run in duplicate. At the conclusion of each experiment, the sand in the columns was dried and kept for solid phase analysis. At most solution conditions, control experiments were performed with empty columns.
13.2.3. Analytical Methods Dissolved lead was analyzed by ICP-OES (Varian Liberty II) or inductively coupled plasma-mass spectrometry (ICP-MS, Agilent Technologies 7500ce). For lead, the detection limit of ICP-OES was 0.01 ppm and that of ICP-MS was 0.01 ppb. The acid matrix used for ICP-OES was 2% HNO3 and for ICP-MS it was 1% HNO3. The calibration standards for ICP analysis were trace ICP/ICP-MS grade (Fisher Scientific). Indium was used as an internal standard for ICP-MS analysis. Dissolved phosphate was measured by the ascorbic acid method (American Public Health Association, 1999). In this method ammonium molybdate and potassium antimonyl tartrate react in acid medium with orthophosphate to form phosphomolybdic acid, which can be reduced by
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Table 13.3: Composition of Solutions Used to Generate Saturation Indices at Different pH Values (Concentrations in M). pH
4 5 6 7
SI ¼ 3
SI ¼ 1
SI ¼ +1
SI ¼ +5
[Pb]init
[P]init
[Cl]init
[Pb]init
[P]init
[Cl]init
[Pb]init
[P]init
[Cl]init
[Pb]init
[P]init
[Cl]init
5 107 5 107 5 107 5 107
1.2 105 3.9 107 1.3 108 7.6 1010
4.1 106 1.3 107 4.4 109 2.5 1010
5 107 5 107 5 107 5 107
3.9 105 1.2 106 4.2 108 2.4 109
1.3 105 4.1 107 1.4 108 8.0 1010
5 107 5 107 5 107 5 107
1.2 104 3.9 106 1.3 107 7.6 109
4.1 105 1.3 106 4.4 108 2.5 109
5 107 5 107 5 107 5 107
1.2 103 3.9 105 1.3 106 7.6 108
4.1 104 1.3 105 4.4 107 2.5 108
Influence of Phosphate on Adsorption and Surface Precipitation
359
ascorbic acid to molybdenum blue. The absorbance of the reacted sample solution was measured spectrophotometrically (Perkin-Elmer Lambda 2S) at 880 nm using a 10 cm pathlength quartz cuvette. The detection limit of phosphorus by this method was 1 ppb. Solution pH was measured with a glass pH electrode and a pH meter (Accumet Research AR25). 13.2.4. Geochemical Equilibrium Modeling The software program FITEQL 4.0 was used to determine chemical equilibrium constants of surface complexation reactions. The program uses a non-linear least-squares optimization method to determine parameters of equilibrium models that provide the best fits to experimental datasets (Herbelin and Westall, 1999). The program contains some of the most common models for accounting for electrostatic effects on surface complexation, including the constant capacitance, diffuse layer, and triple layer models. The constant capacitance model was used in this study to account for the electrostatic interactions between the charged solid surface and ionic adsorbates. Although the constant capacitance model is often used for systems of high ionic strength (Stumm, 1992), it was selected in this study because of its limited number of parameters and the availability of previously published single solute adsorption studies for comparison. Two types of adsorption sites were considered in the SCM: the surface hydroxyls associated with goethite (RFeOH) and with quartz (RSiOH). If goethite coating covers most of the surface of the quartz sand or if RSiOH sites do not have a significant effect on lead and phosphate adsorption, then only one type of adsorption site (RFeOH) would need to be considered. The equilibrium constants for the surface acid–base reactions are taken from the published literature for RSiOH sites (Sverjensky and Sahai, 1996) and RFeOH sites (Gunneriusson et al., 1994) (reactions S1–2 and S10–11 in Table 13.4). Surface site densities for RSiOH and RFeOH were set at the value of 2.3 sites/nm2 suggested by Davis and Kent (1990) for oxide minerals. Experimental data were also used in FITEQL to optimize surface site densities, but if optimized site densities were not substantially different from 2.3 sites/nm2, then the value of 2.3 sites/nm2 was used. The ability of equilibrium constants for single sorbate surface complexation reactions from references to simulate our experimental data was evaluated. If the constants from the literature did not provide a good fit, then they were further optimized in FITEQL. This optimization proceeded in two separate steps. First, the equilibrium constants for lead and phosphate surface complexation at
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L. Xie and D. E. Giammar
Table 13.4: Reactions Considered in Aqueous Lead Phosphate System.
P1 P2 P3 P4 S1 S2 S3 S4 S5 S6 S7 S8 S9 S10 S11 S12 S13
Reaction
logKa
Ref.b
Pb5(PO4)3Cl(s)Q5Pb2++3PO3 4 +Cl 2+ 3 Pb5(PO4)3OH(s)Q5Pb +3PO4 +OH + PbHPO4(s)QPb2++PO3 4 +H 2+ – Pb(OH)2(s)QPb +2OH + RFeOH+ 2 2RFeOH+H + RFeOH2RFeO +H RFeOH+Pb2+2RFeOPb++H+ RFeOH+Pb2+2RFeOHPb2+ + RFeOH+PO3 4 +3H 2RFePO4H2+H2O 3 RFeOH+PO4 +2H+2RFePO4H+H2O + 2 RFeOH+PO3 4 +H 2RFePO4 +H2O 2+ 3 + RFeOH+Pb +PO4 +H 2RFePO4Pb+H2O + RFeOH+Pb2++PO432RFeOPbPO 4 +H + + RSiOH2 2RSiOH+H RSiOH2RSiO+H+ RSiOH+Pb2+2RSiOPb++H+ + 2– RSiOH+PO3 4 +H 2RFePO4 +H2O
84.43 76.8 23.81 19.85 7.47 9.51 0.92 9.47 31.13 25.95 18.45
1 1 1 1 2 2 6 6 3 6 6
RFeOH site density ¼ 2.3 sites/nm2 RSiOH site density ¼ 2.3 sites/nm2 Capacitance ¼ 1.28 F/m2
c c
2.4 8.4 1.02
4 4 6
c
5 5 2
a
logK values for surface reactions are for the intrinsic equilibrium constants (Kint). 1: Schecher and McAvoy (1998); 2: Gunneriusson et al. (1994); 3: Nilsson et al. (1992); 4: Sverjensky and Sahai (1996); 5: Davis and Kent (1990); 6: This study. c Reactions were not needed for this study. b
RSiOH sites (reactions S12 and S13 in Table 13.4) were optimized using the datasets for lead and phosphate adsorption to uncoated sand. These values were then held constant as the equilibrium constants for complexation to RFeOH sites (reactions S3–S7 in Table 13.4) were optimized with the datasets for adsorption to goethite-coated sand. Finally, the site (RFeOH and RSiOH) densities and equilibrium constants from single sorbate experiments were combined to assess whether or not they effectively simulate the dual sorbate experimental data. If the combination of single adsorbate reactions did not sufficiently predict dual adsorbate adsorption, then ternary surface complexation reactions (reactions S8 or S9 in Table 13.4) could be added to the model and their equilibrium constants optimized in FITEQL. Precipitation reactions of cholorpyromorphite and lead hydroxide (reactions P1–P4 in Table 13.4) were also included to evaluate their impact on the simulation of lead uptake.
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The software program MINEQL+ (Schecher and McAvoy, 1998) was used to calculate equilibrium speciation based on fixed reactions, constants, and total concentrations. Equilibrium constants for all aqueous phase reactions were those in the MINEQL+ database (Schecher and McAvoy, 1998).
13.3. Results and Discussion 13.3.1. Single Sorbate Adsorption 13.3.1.1. Effect of Goethite Coating on Adsorption The amount of adsorbed lead was always higher on the goethite-coated sand than on the uncoated sand (Fig. 13.2a). However, the iron oxide content was apparently not high enough for goethite to completely dominate adsorption of lead. As seen in SEM images, the sand surface was not completely covered by goethite particles. The quartz surface has sufficient adsorption capacity and affinity for lead that it must be considered in subsequent modeling. The amount of adsorbed phosphate was always higher on the goethitecoated sand than on the uncoated sand (Fig. 13.2b), which indicated that the goethite coating dominated phosphate adsorption. Phosphate adsorption to quartz was so insignificant that adsorption to the quartz was neglected in modeling. Both lead and phosphate adsorb to the goethite-coated sand with patterns typical of the Langmuir isotherm. Because of more favorable adsorption of phosphate than of lead at pH 6, the initial slope of the isotherm is steeper for phosphate than it is for lead.
13.3.1.2. Effect of pH on Lead Adsorption The adsorption of lead on goethite-coated sand was highly pH dependent (Fig. 13.3), increasing from minimal adsorption at pH 4 to nearly complete adsorption at pH 7. Similar pH edges have been observed for lead adsorption to pure goethite (Hayes and Leckie, 1986; Gunneriusson et al., 1994) and are typical of metal cation adsorption to iron oxides (Benjamin, 2002). The SCM, which was used to fit the experimental data of lead adsorption to goethite-coated sand, included two types of adsorption sites, quartz surface hydroxyl (RSiOH) and goethite surface hydroxyl (RFeOH) groups, which were each described by two surface acid–base reactions (S1, S2, S10,
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Figure 13.2: Adsorption of (a) Lead and (b) Phosphate at pH 6, Solid:Liquid Ratio ¼ 50:1 g/l, and 24 h Contact Time. The Ionic Strength was Fixed by NaNO3 at 103 M and 103 M MES Buffer for Lead Adsorption and 102 M and 102 M MES for Phosphate Adsorption.
and S11 in Table 13.4). In the constant capacitance model, a capacitance of 1.28 F/m2 was used based on previous studies (Gunneriusson et al., 1994). The equilibrium constant for lead adsorption to RSiOH groups, with the site density fixed at 2.3 sites/nm2, was optimized by FITEQL as 101.02 using the dataset for lead adsorption to uncoated sand. The quality of the model fit is characterized by a weighted sum of squares divided by the degrees of freedom (WSOS/DF) value of 0.95. The constant for reaction S12 in Table 13.4 was then held constant in modeling lead adsorption to goethitecoated sand. A site density of 2.3 sites/nm2 was also used for RFeOH sites.
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Figure 13.3: Adsorption Edges (a) and Surface Speciation (b) for Lead Adsorption to Goethite-Coated Sand as a Function of pH. Data are Shown with Points and Surface Complexation Modeling Simulations are Shown by the Solid Line. Reaction Conditions: I ¼ 1.2–1.8 103 M NaNO3, Contact Time ¼ 24 h, Solid:Liquid Ratio ¼ 50:1 g/l.
The optimal equilibrium constants for lead adsorption on goethite were determined to be 100.92 and 109.47 for reactions S3 and S4, respectively, with a WSOS/DF value of 2.7. The equilibrium constants are similar to those determined by Gunneriusson and coworkers (1994) for lead adsorption to pure goethite (100.17 and 108.20). Initial modeling was attempted using a single adsorption reaction to goethite, but the model with one reaction could not provide a good fit to the experimental data over the full range of total lead loading. Inclusion of a second reaction significantly improved the quality of the fit over a broader range of total concentration, although the model
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still overpredicts adsorption for the highest lead concentration. The optimization sought the best fit for 5 107 M total lead, because this concentration was also studied in the presence of phosphate. The total site density of the goethite-coated quartz sand was set at 2.3 sites/nm2 for both goethite and quartz surfaces. This value was recommended by Davis and Kent (1990) for mineral oxide surfaces as one for which the results of multiple studies can be compared. When the goethite site density for phosphate adsorption to goethite-coated sand was simultaneously optimized with the equilibrium constant for reaction S7 in Table 13.4, a value of 2.2 sites/nm2 was determined (see next section), and so the value of 2.3 sites/nm2 was used. The site density for goethite is within the range of values used in previous work on lead adsorption to pure goethite of 1.7 sites/nm2 (Gunneriusson et al., 1994) and 7 sites/nm2 (Hayes and Leckie, 1986). The site density of 2.3 sites/nm2 used for quartz is lower than values of 5.9 sites/nm2 (Yee and Fein, 2003) and 5 sites/nm2 (Chen et al., 2006) used in previous studies. Without an independent measurement of site density, the determinations of the site density of surface groups and the equilibrium binding constants are not entirely independent in model derivation based on adsorption data sets. The use of a consistent site density is beneficial for comparison of equilibrium constants from different studies and is important for the development of a self-consistent thermodynamic database (Davis and Kent, 1990). Figure 13.3b shows the simulated surface speciation of lead adsorption to goethite-coated sand. The binding of lead to goethite is stronger than that to quartz and RFeOHPb2+ is the dominant surface species over the experimental pH range. At higher pH (>7.5), RFeOHPb2+ is surpassed by RFeOPb+ as the dominant species. The species RSiOPb+ is never dominant in this model, although its contribution increases with increasing pH.
13.3.1.3. Effect of pH on Phosphate Adsorption The adsorption of phosphate on goethite-coated sand was pH dependent (Fig. 13.4). The phosphate adsorption density decreased with increasing pH. These data were similar to previous observations of phosphate adsorption (Cheng et al., 2004). When the phosphate concentration was low ([P]tot ¼ 5 106 M), complete adsorption on goethite-coated sand was approached when the pH was less than 7. For a total phosphate concentration of 5 105 M, adsorbed phosphate was never more than 35% of the total. Adsorption of phosphate to uncoated quartz sand was insignificant compared with that to goethite-coated sand, therefore, the model used for
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Figure 13.4: Adsorption Edges (a) and Surface Speciation (b) for Phosphate on Goethite-Coated Sand as a Function of pH. Data are Shown with Points and Surface Complexation Modeling Simulations are Shown by the Solid Line. Reaction Conditions: I ¼ 1.2–1.8 102 M NaNO3, Contact Time ¼ 24 h, Ptot ¼ 107 to 104 M, Solid:Liquid Ratio ¼ 50:1 g/l. phosphate adsorption to goethite-coated sand only considered adsorption to sites on goethite (RFeOH). The equilibrium constants of acid–base goethite surface reactions, the site density, and the capacitance were the same as those in the SCM for lead adsorption. Three reactions (S5, S6 and S7 in
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Table 13.4) were included for phosphate adsorption, consistent with previous studies (Nilsson et al., 1992; Gao and Mucci, 2003). Initially the equilibrium constants for reactions S5 and S6 in Table 13.4 were set at values from a previous study of phosphate adsorption to goethite (Nilsson et al., 1992), and the constant for reaction S7 in Table 13.4 was determined by finding the value that provided the optimal agreement between the experimental data and the model simulation (Table 13.4). The optimized logK for reaction S7 in Table 13.4 was 18.45 and the goodness of fit was characterized by a WSOS/DF of 6.5. Because adsorption was already at the maximum for a given total phosphate concentration below pH 7, the quality of the fit was insensitive to the constants for reactions S5 and S6 in Table 13.4; however, optimization of the effect of phosphate adsorption on surface charge for dual sorbate experiments (discussed in the section ‘‘Co-Sorption of Lead and Phosphate in Combined Systems’’) required refinement of the constant for reaction S6 in Table 13.4 to a logK of 25.95. The values of 25.95 and 18.45 for reactions S6 and S7 in Table 13.4, respectively, are lower than the values of 26.38 and 20.61 used by Nilsson and coworkers. is The species RFePO4H is dominant at lower pH and RFePO2 4 increase dominant at higher pH (Fig. 13.4b). Concentrations of RFePO2 4 and those of RFePO4H decrease with increasing pH. The species RFePO4H2 is not significant over the experimental pH range. At higher is more total phosphate concentrations, the formation of RFePO2 4 significant (modeling data not shown). 13.3.2. Co-Sorption of Lead and Phosphate in Combined Systems When lead and phosphate were combined in the same systems, the presence of phosphate increased lead removal for all experiments. The effect was greatest at low pH. At pH 4, lead uptake increased from 6% to 40–54% when the initial conditions of the system were undersaturated with respect to pyromorphite and to 60–86% in systems with initial supersaturation. At higher pH (6 or 7), the presence of phosphate did not have as significant of an impact since lead removal was already approaching 100% in the absence of phosphate. Phosphate sorption in the dual sorbate experiments was measured and was generally consistent with the SCM. Three scenarios were considered as models for the enhancement of lead uptake by phosphate: (1) direct combination of binary surface complexation reactions; (2) combination of binary surface complexation reactions with an additional ternary surface complexation reaction (Table 13.4, reaction S8 or S9); (3) combination of binary surface complexation reactions with potential
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Figure 13.5: Lead Sorption on Goethite-Coated Sand as a Function of pH and Solution Saturation with Respect to Pyromorphite. Fixed Reaction Conditions: I ¼ 1.2–1.8 102 M NaNO3, Contact Time ¼ 24 h, Solid:Liquid Ratio ¼ 50:1 g/l, Pbtot ¼ 5 107 M. Simulations Involve (a) Direct Combination of Binary Adsorption Reactions and (b) Binary Adsorption and Precipitation Reactions. precipitation of pyromorphite and lead hydroxide (Table 13.4, reactions P1 and P4). Figure 13.5a shows the simulation from directly combining surface complexation reactions from lead-only and phosphate-only systems. The model does predict that phosphate addition increases lead adsorption. The increase is caused by the lowering of the surface charge and surface potential (c) from the formation of RFePO4H and RFePO2 4 species, which consequently
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Figure 13.6: Effect of Total Phosphate Loading on Goethite-Coated Sand Surface Charge at pH 4. Simulation Conditions: I ¼ 1.2–1.8 102 M NaNO3, Solid:Liquid Ratio ¼ 50:1 g/L, Pbtot ¼ 5 107 M. Total Phosphate Concentrations Varied from 109 to 103 M (Logarithmic Scale). increases lead adsorption as a result of a more favorable electrostatic component for adsorption (Fig. 13.6). The apparent equilibrium constant for reaction S3 in Table 13.4 (Eq. (13.10)) increases with decreasing surface potential (c) at pH 4 because the intrinsic constant remains fixed (Fig. 13.6). K app ¼ K int eF c=RT ¼
f FeOPbþ gfHþ g fFeOHgfPb2þ g
(13.10)
These simulations (Fig. 13.5a) show a good fit to experimental data at pH 4 and 5. Successfully fitting the lead sorption results at pH 4 did require model refinements beyond those initially attempted. Although only optimization of the constant for reaction S7 in Table 13.4 was necessary to model phosphate adsorption, the resulting effect of phosphate adsorption on surface charge overestimated the increase in lead sorption in the presence of phosphate. Optimization of both reactions S6 and S7 in Table 13.4 provided an equally good fit of phosphate sorption but then provided an improved fit for lead sorption in the dual sorbate experiments. The model predicts decreasing lead adsorption at pH 4.5–5 because the total phosphate concentrations at pH 4.5–5 are less than at pH 4 in order to maintain the same initial saturation indices with respect to pyromorphite (Table 13.3), although this effect remains to be tested experimentally.
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The inclusion of reactions for pyromorphite and lead hydroxide precipitation (reactions P1 and P4 in Table 13.4) only predict higher uptake at the initial saturation index of +1 at pH less than 3.5 and at the initial SI of +5 at pH o 4.5. Pyromorphite was the solid predicted to form at these conditions. At lower initial saturation conditions and higher pH, adsorption was the dominant mechanism for lead uptake. Adsorption reduced dissolved lead and phosphate to concentrations that would no longer be supersaturated. The model simulations are based on equilibrium conditions, but in the actual experimental systems the sorption mechanism may be determined by the relative rates of adsorption and precipitation. The assumption of equilibrium also suggests that there would not be an effect of the order of sorbate addition in dual sorbate experiments; however, this is an effect that should be examined in future work. Inclusion of a ternary surface complexation reaction as a ligand-bridging surface complex (reaction S8 in Table 13.4) did not improve the model fit to the experimental data, nor were other reaction stoichiometries (e.g., metalbridging complexes) helpful for improving the fit. No values for a ternary surface complex improved the model fit. The combination of single-sorbate SCMs already fit the data or slightly overestimated lead uptake, and the inclusion of ternary surface complexes would only increase the predicted uptake. Previous studies of electrophoretic mobility and X-ray absorption spectroscopy show that adsorbed phosphate acts as a reactant to form lead phosphate surface phases that may be highly dispersed on goethite (Weesner and Bleam, 1998). In previous experiments by Ler and Stanforth (2003), the reaction stoichiometries of lead adsorption on phosphated goethite suggested that pyromorphite (Pb5(PO4)3Cl) may precipitate on the surface at conditions that are undersaturated with respect to pyromorphite solubility. The surface precipitation of lead phosphate on the goethite surface is most likely by formation of a new surface phase with the structure of pyromorphite. Analogous research on the co-sorption of zinc and arsenate, a good analog for phosphate, on goethite found that the presence of zinc increased arsenate adsorption and vice versa. The effect was most pronounced at high surface loadings and EXAFS data were used to determine that arsenate and zinc were present in zinc arsenate surface precipitates even at conditions that were undersaturated with respect to zinc arsenate solid phases (Gra¨fe et al., 2004; Gra¨fe and Sparks, 2005). Although these experiments showed no evidence of pyromorphite by SEM and XRD, precipitation of lead phosphate minerals cannot be ruled out. Since the maximum loading of lead on the solid is only 2 mg/kg, the formation of lead phosphate minerals may not be detectable by these
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techniques. The low lead loading would even be challenging to detect by element-specific spectroscopic techniques such as EXAFS. Further investigation with goethite-coated sand on the nature of surface precipitation can focus on conditions of higher initial saturation index and lower pH. The low mass loading of lead on the solid is the result of working with goethite-coated sand, which was selected because of its potential utility as a model geomedium, and valuable information on mechanisms that will apply to this material can be gained from the numerous spectroscopic studies previously conducted with pure goethite.
13.4. Summary Goethite-coated sand can serve as a model porous medium to study the fate of lead in natural soil systems. Lead adsorption to goethite-coated sand can be simulated by a SCM that accounts for adsorption to goethite and quartz surface sites, and phosphate can be modeled using just goethite surface sites. The site densities, reaction stoichiometries, and equilibrium constants for goethite-coated sand are similar to those of pure goethite. The chemical equilibrium model developed in this study can be incorporated into reactive transport models for lead and phosphate transport through sand columns and ultimately to transport in soil and groundwater systems. Over a broad range of solution compositions, phosphate addition enhanced lead removal. Lead sorption can be predicted by the combination of lead and phosphate surface complexation reactions. Ternary surface complexation is not necessary for uptake simulation. The SCM suggests that the dominant sorption mechanism is adsorption, even when the system is initially supersaturated with respect to the lead phosphate mineral pyromorphite. Precipitation mechanisms will become significant at lower pH and higher supersaturation conditions. Since the loading of lead for this study is lower than the detection limits of most techniques for solid phase investigation, the possibility of pyromorphite formation cannot be ruled out based only on macroscopic sorption data.
ACKNOWLEDGMENTS Partial funding for this study was provided by the National Science Foundation (BES #0546219). Liyun Xie has been supported by fellowships through the Washington University School of Engineering and Applied Science, the Cecil Lue-Hing Scholarship, and the Lilia Abron Scholarship.
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The authors acknowledge Claire Farnsworth’s contributions to lead adsorption experiments and Yu Wang’s to phosphate adsorption experiments. The insightful comments of Doug Kent and three anonymous reviewers assisted in the revision of this chapter.
REFERENCES American Public Health Association A. W. W. A. (1999). Water Environment Federation. Standard Methods for the Examination of Water and Wastewater. Arai, Y., & Sparks, D. L. (2001). ATR-FTIR spectroscopic investigation on phosphate adsorption mechanisms at the ferrihydrite-water interface. J. Colloid Interface Sci., 241 (2), 317–326. Bargar, J. R., Brown, G. E., & Parks, G. A. (1997). Surface complexation of Pb(II) at oxide-water interfaces. 2. XAFS and bond-valence determination of mononuclear Pb(II) sorption products and surface functional groups on iron oxides. Geochim. Cosmochim. Acta, 61 (13), 2639–2652. Benjamin, M. M. (2002). Water Chemistry. McGraw-Hill, New York, NY. Cao, X., Ma, L. Q., Chen, M., Singh, S. P., & Harris, W. G. (2002). Impacts of phosphate amendments on lead biogeochemistry at a contaminated site. Environ. Sci. Technol., 36 (24), 5296–5304. Chen, C. C., Coleman, M. L., & Katz, L. E. (2006). Bridging the gap between macroscopic and spectroscopic studies of metal ion sorption at the oxide/water interface: Sr(II), Co(II), and Pb(II) sorption to quartz. Environ. Sci. Technol., 40 (1), 142–148. Cheng, T., Barnett, M. O., Roden, E. E., & Zhuang, J. (2004). Effects of phosphate on uranium(VI) adsorption to goethite-coated sand. Environ. Sci. Technol., 39 (22), 6059–6065. Clark, S. B., Johnson, W. H., Malek, M. A., Serkiz, S. M., & Hinton, T. G. (1996). A comparison of sequential extraction techniques to estimate geochemical controls on the mobility of fission product, actinide, and heavy metal contaminants in soils. Radiochim. Acta, 74, 173–179. Coston, J. A., Fuller, C. C., & Davis, J. A. (1995). Pb2+ and Zn2+ adsorption by a natural aluminum- and iron-bearing surface coating on an aquifer sand. Geochim. Cosmochim. Acta, 59 (17), 3535–3547. Davis, J. A., & Kent, D. B. (1990). Surface complexation modeling in aqueous geochemistry. In: M. F. Hochella, & A. F. White (Eds). Mineral-Water Interface Geochemistry, Vol. 23. Mineralogical Society of America, Washington, DC, pp. 177–260. Dideriksen, K., & Stipp, S. L. S. (2003). The adsorption of glyphosate and phosphate to goethite: A molecular-scale atomic force microscopy study. Geochim. Cosmochim. Acta, 67 (18), 3313–3327. Dzombak, D. A., & Morel, F. M. M. (1990). Surface Complexation Modeling. Wiley-Interscience, New York, NY.
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Elzinga, E. J., & Sparks, D. L. (2002). X-ray absorption spectroscopy study of the effects of pH and ionic strength on Pb(II) sorption to amorphous silica. Environ. Sci. Technol., 36 (20), 4352–4357. Gabriel, U., Gaudet, J.-P., Spadini, L., & Charlet, L. (1998). Reactive transport of uranyl in a goethite column: An experimental and modelling study. Chem. Geol., 151, 107–128. Gao, Y., & Mucci, A. (2003). Individual and competitive adsorption of phosphate and arsenate on goethite in artificial seawater. Chem. Geol., 199 (1–2), 91–109. Good, N. E., Winget, G. D., Winter, W., Connolly, T. N., Izawa, S., & Singh, R. M. M. (1966). Hydrogen ion buffers for biological research. Biochemistry, 5 (2), 467–477. Gra¨fe, M., Nachtegaal, M., & Sparks, D. L. (2004). Formation of metal-arsenate precipitates at the goethite-water interface. Environ. Sci. Technol., 38 (24), 6561–6570. Gra¨fe, M., & Sparks, D. L. (2005). Kinetics of zinc and arsenate co-sorption at the goethite-water interface. Geochim. Cosmochim. Acta, 69 (19), 4573–4595. Gunneriusson, L., Lovgren, L., & Sjoberg, S. (1994). Complexation of Pb(Ii) at the geothite (alpha-FeOOH) water interface – the influence of chloride. Geochim. Cosmochim. Acta, 58 (22), 4973–4983. Hayes, K. F., & Leckie, J. O. (1986). Mechanism of lead ion adsorption at the goethite-water interface. In: J. A. Davis, & K. F. Hayes (Eds). Geochemical Processes at Mineral Surfaces, Vol. 323. American Chemical Society, Washington, DC, pp. 114–141. Herbelin, A. L., & Westall, J. C. (1999). FITEQL – A Computer Program for Determination of Chemical Equilibrium Constants from Experimental Data. Oregon State University, OR. Hering, J. G., & Kraemer, S. (1994). Kinetics of complexation reactions at surfaces and in solution: Implications for enhanced radionuclide migration. Radiochim. Acta, 66/67, 63–71. Hettiarachchi, G. M., Pierzynski, G. M., & Ransom, M. D. (2000). In situ stabilization of soil lead using phosphorus and manganese oxide. Environ. Sci. Technol., 34 (21), 4614–4619. Ler, A., & Stanforth, R. (2003). Evidence for surface precipitation of phosphate on goethite. Environ. Sci. Technol., 37 (12), 2694–2700. Ma, Q. Y., Logan, T. J., & Traina, S. J. (1995). Lead immobilization from aqueoussolutions and contaminated soils using phosphate rocks. Environ. Sci. Technol., 29 (4), 1118–1126. Nilsson, N., Lovgren, L., & Sjoberg, S. (1992). Phosphate complexation at the surface of goethite. Chem. Speciation Bioavailability, 4 (4), 121–130. Nriagu, J. O. (1973). Lead orthophosphates – II. Stability of chloropyromorphite at 251C. Geochim. Cosmochim. Acta, 37, 367–377. Roe, A. L., Hayes, K. F., Chisholmbrause, C., Brown, G. E., Parks, G. A., Hodgson, K. O., & Leckie, J. O. (1991). In situ X-ray absorption study of lead-ion surface complexes at the goethite water interface. Langmuir, 7 (2), 367–373.
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Ruby, M. V., Davis, A., & Nicholson, A. (1994). In-situ formation of lead phosphates in soils as a method to immobilize lead. Environ. Sci. Technol., 28 (4), 646–654. Ryan, J. A., Scheckel, K. G., Berti, W. R., Brown, S. L., Casteel, S. W., Chaney, R. L., Hallfrisch, J., Doolan, M., Grevatt, P., Maddaloni, M., & Mosby, D. (2004). Reducing children’s risk from lead in soil. Environ. Sci. Technol., 38 (1), 18A–24A. Schecher, W. D., & McAvoy, D. C. (1998). MINEQL+: A Chemical Equilibrium Modeling System. Version 4.5. Environmental Research Software. Scheidegger, A., Borkovec, M., & Sticher, H. (1993). Coating of silica sand with goethite: Preparation and analytical identification. Geoderma, 58, 43–65. Schwertmann, U., & Cornell, R. M. (2000). Iron Oxides in the Laboratory. Wiley-VCH, New York, NY. Soares, H. M. V. M., Conde, P. C. F. L., Almeida, A. A. N., & Vasconcelos, M. T. S. D. (1999). Evaluation of n-substituted aminosulfonic acid pH buffers with a morpholinic ring for cadmium and lead speciation studies by electroanalytical techniques. Anal. Chim. Acta, 394, 325–335. Sposito, G. (1986). Distinguishing adsorption from surface precipitation. In: J. A. Davis, & K. F. Hayes (Eds). Geochemical Processes at Mineral Surfaces, Vol. 323. American Chemical Society, Washington, DC, pp. 217–228. Stumm, W. (1992). Chemistry of the Solid-Water Interface. Wiley, New York, NY. Sverjensky, D. A., & Sahai, N. (1996). Theoretical prediction of single-site surfaceprotonation equilibrium constants for oxides and silicates in water. Geochim. Cosmochim. Acta, 60 (20), 3773–3797. Trivedi, P., Dyer, J. A., & Sparks, D. L. (2003). Lead sorption onto ferrihydrite. 1. A macroscopic and spectroscopic assessment. Environ. Sci. Technol., 37 (5), 908–914. Weesner, F. J., & Bleam, W. F. (1998). Binding characteristics of Pb2+ on anionmodified and pristine hydrous oxide surfaces studied by electrophoretic mobility and X-ray absorption spectroscopy. J. Colloid Interface Sci., 205 (2), 380–389. Yang, J., Mosby, D. E., Casteel, S. W., & Blanchar, R. W. (2001). Lead immobilization using phosphoric acid in a smelter-contaminated urban soil. Environ. Sci. Technol., 35 (17), 3553–3559. Yee, N., & Fein, J. B. (2003). Quantifying metal adsorption onto bacteria mixtures: A test and application of the surface complexation model. Geomicrobiol. J., 20 (1), 43–60. Zhang, P. C., Ryan, J. A., & Bryndzia, L. T. (1997). Pyromorphite formation from goethite adsorbed lead. Environ. Sci. Technol., 31 (9), 2673–2678.
Chapter 13
Influence of Phosphate on Adsorption and Surface Precipitation of Lead on Iron Oxide Surfaces Liyun Xie and Daniel E. Giammar Department of Energy, Environmental & Chemical Engineering, Washington University in St. Louis, One Brookings Dr. St. Louis, MO 63130, USA
ABSTRACT Lead and phosphate sorption on goethite-coated and uncoated quartz sand was measured experimentally and modeled within a reaction-based framework. Single sorbate batch experiments and experiments with lead and phosphate present together were conducted. Adsorption of lead and phosphate to goethite-coated sand was dominated by adsorption to the goethite coating. A surface complexation model adapted from models for pure goethite successfully simulated lead and phosphate adsorption to goethite-coated sand over a broad range of pH and total sorbate concentrations; the inclusion of a surface complexation reaction for lead adsorption to the quartz surface was necessary to improve the model fit. Lead sorption on goethite-coated sand was enhanced by the presence of phosphate. The effect was most pronounced at low pH. The enhanced lead uptake was predicted by the combination of the single sorbate surface complexation models. The adsorption of phosphate at low pH decreased the surface charge and potential, which increases the extent of lead adsorption. No ternary surface complexes were needed to model the dual sorbate results. When a reaction for the precipitation of chloropyromorphite (Pb5(PO4)3Cl) was included, the model predicted precipitation only at the lowest pH and highest phosphate loading studied. Over most experimental conditions, including conditions that were initially supersaturated with respect to chloropyromorphite, the equilibrium model predicted that adsorption was the dominant mechanism of lead sorption; however, the actual mechanisms may be controlled by the relative rates of precipitation and adsorption reactions.
Corresponding author. Tel.: +314-935-6849; Fax: +314-935-5464;
E-mail: [email protected] (D.E. Giammar).
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13.1. Introduction The fate and transport of lead in soil and groundwater is strongly controlled by chemical reactions at the mineral–water interface. Surface reactions include adsorption and precipitation. Goethite (a-FeOOH), a common iron(III) oxyhydroxide mineral in natural systems, and other iron(III) oxides and oxyhydroxides are important environmental sorbents for heavy metals due to their high specific surface areas and reactive surfaces (Dzombak and Morel, 1990; Coston et al., 1995; Schwertmann and Cornell, 2000). Iron oxide coated sands prepared in the laboratory have previously been used as model materials for real soils and sediments that have iron oxide coatings (Gabriel et al., 1998; Cheng et al., 2004). Lead adsorbs to goethite through the formation of inner-sphere surface complexes with surface hydroxyl groups (RFeOH) as determined by extended X-ray adsorption fine-structure spectroscopy (EXAFS) (Roe et al., 1991; Bargar et al., 1997). Lead also adsorbs via inner-sphere complexation as a mixture of monodentate and bidentate complexes on ferrihydrite (Fe5HO8 4H2O) (Trivedi et al., 2003) and as mononuclear bidentate complexes on hematite (a-Fe2O3) (Bargar et al., 1997). Evidence for inner-sphere complexation is also provided by a lack of ionic strength effects on adsorption. The ionic strength of the bulk solution has less impact on inner-sphere complexation than on outer-sphere complexation (Hayes and Leckie, 1986). Adsorption of lead to goethite has been successfully interpreted through the inclusion of reaction 1 in a surface complexation model (SCM) (Hayes and Leckie, 1986). FeOH þ Pb2þ ¼ FeOPbþ þ Hþ
(13.1)
Other SCMs have included reaction 1 as well as two other reactions (2 and 3) (Gunneriusson et al., 1994). FeOH þ Pb2þ ¼ FeOHPb2þ
(13.2)
FeOH þ Pb2þ þ H2 O ¼ FeOPbOH þ 2Hþ
(13.3)
Surface complexation modeling considers the adsorption of inorganic species to specific reactive surface sites and includes terms to account for both the chemical and electrostatic energetics of adsorption (Dzombak and Morel, 1990). The model of Hayes and Leckie (1986) successfully predicted the pH dependence of lead adsorption. Although metal adsorption to quartz is not usually significant in soil and groundwater relative to adsorption to iron oxyhydroxides and clays, metal
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cations can still adsorb at the quartz surface. X-ray absorption spectroscopy investigations found inner-sphere lead surface complexes formed on quartz (Chen et al., 2006) and amorphous silica (Elzinga and Sparks, 2002). Lead sorption at the iron oxide–water interface can be enhanced by the presence of phosphate. Enhancement can occur as the result of: (a) changes to surface charge that make lead adsorption more favorable, (b) formation of ternary lead–phosphate–iron oxide surface complexes, (c) precipitation of lead phosphate solids, and (d) surface alteration from the formation of an iron(III) phosphate surface precipitate. Phosphate addition to leadcontaminated soils has been proposed as a means of reducing lead mobility and bioavailability (Ruby et al., 1994; Ma et al., 1995; Hettiarachchi et al., 2000; Yang et al., 2001; Cao et al., 2002; Ryan et al., 2004). Although remediation strategies usually suggest using phosphate addition to precipitate lead phosphate solids, enhancement of lead binding to iron oxyhydroxides may also increase lead sorption even without lead phosphate precipitation. Phosphate adsorbs to goethite through inner-sphere complex formation as observed directly with attenuated total reflectance Fourier transform infrared spectroscopy (Arai and Sparks, 2001). Atomic force microscopy (AFM) of phosphate adsorption to the (0 1 0) surface of a goethite single crystal observed phosphate adsorbed in a 1:1 ratio with the singly coordinated hydroxyl groups via monodentate complexes (Dideriksen and Stipp, 2003). A combination of three monodentate surface complexes (reactions 4–6) has been proposed for modeling phosphate adsorption to goethite (Nilsson et al., 1992). þ FeOH þ H2 PO 4 þ H ¼ FePO4 H2 þ H2 O
(13.4)
FeOH þ H2 PO 4 ¼ FePO4 H þ H2 O
(13.5)
2 þ FeOH þ H2 PO 4 ¼ FePO4 þ H þ H2 O
(13.6)
When metal ions (e.g., M2+) and anions (e.g., L) adsorb together on the oxide surface, ternary surface complexes can form by either metal-bridging (reaction 7) or ligand-bridging (reaction 8) reactions (Hering and Kraemer, 1994). SOH þ M2þ þ L ¼ SOM L þ Hþ
(13.7)
SOH þ L þ M2þ þ Hþ ¼ SLM2þ þ H2 O
(13.8)
Ternary surface complexation is an adsorption process with the simultaneous accumulation of metal cations and ligand anions at the solid– water interface without the development of three-dimensional structure
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(Sposito, 1986). This adsorption process is in contrast to the formation of a precipitate at the surface that has longer range three-dimensional structure. Lead and phosphate can also precipitate at the goethite surface, most likely through heterogeneous nucleation of a lead phosphate precipitate on the goethite substrate. Once precipitation is initiated, the number of surface sites for lead or phosphate binding is no longer fixed by the goethite surface area. Surface precipitation is most likely at high sorbate loading. Potential lead-containing solids include lead hydroxide (Pb(OH)2) and the lead phosphates chloropyromorphite (Pb5(PO4)3Cl), hydroxypyromorphite (Pb5(PO4)3OH), and lead hydrogen phosphate (PbHPO4). Ler and Stanforth (2003) observed precipitation of lead and phosphate on goethite in a phase with the stoichiometry of chloropyromorphite or hydroxypyromorphite. The reaction of hydroxyapatite (Ca5(PO4)3OH) and goethite with adsorbed lead resulted in the formation of chloropyromorphite through homogeneous nucleation in the solution and possibly heterogeneous nucleation at the goethite surface (Zhang et al., 1997). In this study we examine the effects of phosphate on lead sorption to goethite-coated sand. The word sorption is used here to describe all processes (adsorption, surface precipitation, co-precipitation) that result in accumulation of lead at the goethite–water interface. The objectives of this project were to determine the utility of goethite-coated sand as a model porous medium with a reactive surface phase, distinguish the different mechanisms through which phosphate can increase lead sorption, and determine the geochemical conditions under which each mechanism is dominant. Once the interfacial chemical reactions on goethite-coated sand have been established, it can be a useful material for evaluating the effects of surface reactions on lead transport. The different mechanisms of lead sorption in the presence of phosphate may result in similar equilibrium partitioning of lead between the solid and dissolved phases, but the rates of transport are likely to be affected by differences in sorption mechanisms.
13.2. Experimental 13.2.1. Materials All chemicals used were certified ACS grade (Fisher Scientific) unless otherwise specified. The concentrated nitric acid was trace metal grade. Pure goethite was synthesized from an alkaline aqueous system according to methods of Schwertmann and Cornell (2000). A 100 ml volume of 1 M
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Fe(NO3)3 was mixed with 180 ml of 5 M KOH in a 2 l polyethylene bottle and then immediately diluted with ultrapure water to 2 l. The bottle was sealed and heated in an oven at 651C for 60 h. An ocher precipitate formed during the heating process. The resulting suspension was washed free of excess dissolved ions, primarily K+ and NO 3 , by dialysis. The dialyzed suspension was freeze dried. The identity of the synthetic goethite (a-FeOOH) was confirmed by X-ray powder diffraction (XRD) (Rigaku Geigerflex D-MAX/A Diffractometer) with Cu–Ka radiation. Solid samples for XRD were ground to powders and loaded into glass sample holders for analysis. Quartz sand (U.S. Silica) with particle sizes of 212–355 mm was soaked in 1 M HCl (Fisher Scientific) overnight to remove surface impurities from the quartz. The sand was rinsed with ultrapure water to remove the acid and then dried at 1051C. Goethite-coated sand was then prepared by agitating a mixture of acid-washed quartz sand and goethite at pH 6.8 and 0.01 M ionic strength (NaNO3) for 24 h. These conditions had previously been determined to be optimal for goethite attachment based on the surface charges of goethite and quartz (Scheidegger et al., 1993). Unattached goethite was removed by repeatedly suspending the sand in 0.01 M NaNO3 solution, allowing the sand to settle, and decanting the supernatant. These steps were repeated until a clear supernatant was achieved. The iron content of the coated sand was determined by extraction in a solution of 0.3 M sodium dithionite, 0.3 M sodium citrate, and 0.2 M sodium bicarbonate. This solution effectively dissolves crystalline Fe(III) oxides and oxyhydroxides by reducing them to soluble Fe(II) species (Clark et al., 1996). The extraction mixture was shaken for 1 h, set overnight at room temperature for complete reaction, and the extract was then analyzed by inductively coupled plasmaoptical emission spectrometry (ICP-OES). The goethite-coated quartz sand contained 0.2% (w/w) iron; the value of 0.005% iron in uncoated sand was reported by the supplier (U.S. Silica). The typical crystalline morphology of goethite and of goethite present on goethite-coated sand was observed with a field emission scanning electron microscope (Hitachi S-4500) (Fig. 13.1). Dry samples were fixed to sample holders with double-sided tape. To improve the SEM images, gold coating was used on some samples to increase electrical conductivity. The quartz sand surface was partially covered by goethite particles but large portions of quartz remained exposed. Table 13.1 summarizes the properties of pure goethite, uncoated quartz sand, and goethite-coated sand. The size of uncoated and goethite-coated sand were the same, but the specific surface area of the coated sand was more than three times that of uncoated sand due to the high surface area goethite coating. Specific surface area (m2/g) was measured by
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Figure 13.1: Scanning Electron Micrographs of Synthetic Goethite (Left) and Goethite-Coated Sand (Right).
Table 13.1: Sorbent Properties.
Size (mm) Iron (wt%) Specific surface area (m2/g)
Goethite
Quartz sand
Goethite-coated sand
1 62.8 33.6
212–355 0.005 0.05
212–355 0.2 0.16
BET-N2 adsorption (Quantachrome Autosorb AS-1). Prior to BET-N2 measurement, wet samples were freeze dried (Labconco, Freezone 4.5) and then loaded into sample cells for analysis. Based on the iron content of the goethite-coated sand and assuming the additional 0.11 m2/g of surface area for the coated sand is from goethite, the goethite in the coating is calculated to have a specific surface area of 35 m2/g, which is very close to the value of 33.6 m2/g measured for the pure goethite.
13.2.2. Recirculating Micro-Column Batch Adsorption Experiments Lead and phosphate sorption reactions were studied by continuously recirculating solution from a 50 ml reservoir through 1.8 ml columns that were filled with 2.5 g of either goethite-coated sand or uncoated sand. Columns with recirculation, instead of stirred batch reactors, were used because goethite detachment from quartz sand was observed when suspensions of goethite-coated sand were agitated on a rotary shaker or with a magnetic stir bar. The recirculation flow rate was controlled at 0.5 ml/min with a peristaltic pump. Samples were collected after 24 h, which allowed sufficient
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residence times of both the column (3 min) and the reservoir (100 min) to allow the entire system to be considered as one well-mixed batch reactor. Lead adsorption to goethite occurs on time scales much faster than those of this experiment (Hayes and Leckie, 1986), and our preliminary experiments demonstrated that 24 h was also sufficient for equilibration. After the 24 h equilibration period, a 20 ml effluent sample was collected for pH and flow-rate measurement and another 10 ml sample was collected for dissolved lead and/or phosphate analysis. For many of the single sorbate experiments the influent reservoir was then refilled with 30 ml of solution to provide a new desired composition and the system was equilibrated for another 24 h until the next sampling time. By using this approach, multiple adsorption datapoints were generated while using a limited amount of solid material and maintaining a completely saturated environment in the columns. Single sorbate (lead-only or phosphate-only) and binary sorbate (lead and phosphate together) experiments were conducted (Table 13.2). Single sorbate experiments were conducted to evaluate the extent to which adsorption to goethite-coated sand could be explained by adsorption to pure goethite. The influent lead and phosphate concentrations were undersaturated with respect to lead hydroxide and iron phosphate precipitates. The range of total lead concentrations was 107–104 M and that of total phosphate concentrations was 107–5 105 M. Lead was added as Pb(NO3)2 (Acros Organic) and phosphate as KH2PO4 (Fisher Scientific). The desired pH of each adsorption experiment was from 4 to 8 and was controlled by MES (Sigma) or HEPES (Acros Organic) buffer solutions adjusted to pH 6–8 by addition of NaOH. These buffers were selected because of their established low affinities for metal complexation (Good et al., 1966; Soares et al., 1999). Buffer concentrations were 1 mM for lead adsorption and 10 mM for phosphate adsorption experiments. The buffer concentrations were selected to be at least an order of magnitude higher than the maximum total adsorbate concentration studied, and phosphate was studied at a higher total concentration than was lead. For control of the ionic strength, solutions also contained 1 mM NaNO3 for lead adsorption experiments and 10 mM with NaNO3 for phosphate adsorption experiments. When combined with the buffer species and associated NaOH was added, the total ionic strength of the solutions was 1.2–1.8 mM for lead experiments and 12–18 mM for phosphate experiments, depending on the pH. To examine the transition between adsorption and surface precipitation, binary sorbate experiments were conducted in which lead and phosphate were both present. The pH in these experiments was kept at 4–7 and the initial lead concentration was fixed at 5 107 M. To provide the desired
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Experiment no.
A-1 A-2 B-1 B-2 C-1 C-2 a
Sorbent
Uncoated sand Goethite-coated sand Uncoated sand Goethite-coated sand Uncoated sand Goethite-coated sand
Solution composition pH
[Pb]init (M)
[P]init (M)
[Cl]init (M)
SICPY,
4–8 4–8
107–104a 107–104a
0 0
0 0
NA NA
4–8 4–8
0 0
107–5 105a 107–5 105a
0 0
NA NA
4–7 4–7
5 107 5 107
7.5 1010–103 7.5 1010–103
2.5 1010–104 2.5 1010–104
3, 1, 1, 5 3, 1, 1, 5
Single columns were used to evaluate adsorption over a broad concentration range at a given pH for lead-only and phosphate-only experiments.
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Table 13.2: Conditions of Experiments Conducted.
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initial saturation index with respect to chloropyromorphite, phosphate concentrations were varied from 7.5 1010 to 1.2 103 M and chloride concentrations from 2.5 1010 to 4 104 M (Table 13.3). These values were selected to approach relevant environmental concentrations while avoiding the use of initial solutions that were at even higher initial conditions of saturation. The initial saturation index (SI) (Eq. (13.9)) was examined at two conditions of undersaturation (3 and 1) and two of supersaturation (+1 and +5). The solubility product (Ksp) of chloropryromorphite is 1084.43 (Nriagu, 1973). 3 fPb2þ g5 fPO3 4 g fCl g SI ¼ log (13.9) K sp At negative values of SI, the precipitation of chloropyromorphite is not expected, and at positive SI values, chloropyromorphite precipitation is thermodynamically favorable. In sorbent-free control experiments, supersaturated conditions could be maintained for times of more than a day without loss of lead or phosphate from solution; however, at longer time scales, chloropyromorphite precipitation was observed for supersaturated conditions. Only the initial saturation indices were set at fixed values, and once lead and phosphate are taken up by the solid, either by adsorption or precipitation, the saturation index will decrease from its initial value. All experiments were run in duplicate. At the conclusion of each experiment, the sand in the columns was dried and kept for solid phase analysis. At most solution conditions, control experiments were performed with empty columns.
13.2.3. Analytical Methods Dissolved lead was analyzed by ICP-OES (Varian Liberty II) or inductively coupled plasma-mass spectrometry (ICP-MS, Agilent Technologies 7500ce). For lead, the detection limit of ICP-OES was 0.01 ppm and that of ICP-MS was 0.01 ppb. The acid matrix used for ICP-OES was 2% HNO3 and for ICP-MS it was 1% HNO3. The calibration standards for ICP analysis were trace ICP/ICP-MS grade (Fisher Scientific). Indium was used as an internal standard for ICP-MS analysis. Dissolved phosphate was measured by the ascorbic acid method (American Public Health Association, 1999). In this method ammonium molybdate and potassium antimonyl tartrate react in acid medium with orthophosphate to form phosphomolybdic acid, which can be reduced by
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Table 13.3: Composition of Solutions Used to Generate Saturation Indices at Different pH Values (Concentrations in M). pH
4 5 6 7
SI ¼ 3
SI ¼ 1
SI ¼ +1
SI ¼ +5
[Pb]init
[P]init
[Cl]init
[Pb]init
[P]init
[Cl]init
[Pb]init
[P]init
[Cl]init
[Pb]init
[P]init
[Cl]init
5 107 5 107 5 107 5 107
1.2 105 3.9 107 1.3 108 7.6 1010
4.1 106 1.3 107 4.4 109 2.5 1010
5 107 5 107 5 107 5 107
3.9 105 1.2 106 4.2 108 2.4 109
1.3 105 4.1 107 1.4 108 8.0 1010
5 107 5 107 5 107 5 107
1.2 104 3.9 106 1.3 107 7.6 109
4.1 105 1.3 106 4.4 108 2.5 109
5 107 5 107 5 107 5 107
1.2 103 3.9 105 1.3 106 7.6 108
4.1 104 1.3 105 4.4 107 2.5 108
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ascorbic acid to molybdenum blue. The absorbance of the reacted sample solution was measured spectrophotometrically (Perkin-Elmer Lambda 2S) at 880 nm using a 10 cm pathlength quartz cuvette. The detection limit of phosphorus by this method was 1 ppb. Solution pH was measured with a glass pH electrode and a pH meter (Accumet Research AR25). 13.2.4. Geochemical Equilibrium Modeling The software program FITEQL 4.0 was used to determine chemical equilibrium constants of surface complexation reactions. The program uses a non-linear least-squares optimization method to determine parameters of equilibrium models that provide the best fits to experimental datasets (Herbelin and Westall, 1999). The program contains some of the most common models for accounting for electrostatic effects on surface complexation, including the constant capacitance, diffuse layer, and triple layer models. The constant capacitance model was used in this study to account for the electrostatic interactions between the charged solid surface and ionic adsorbates. Although the constant capacitance model is often used for systems of high ionic strength (Stumm, 1992), it was selected in this study because of its limited number of parameters and the availability of previously published single solute adsorption studies for comparison. Two types of adsorption sites were considered in the SCM: the surface hydroxyls associated with goethite (RFeOH) and with quartz (RSiOH). If goethite coating covers most of the surface of the quartz sand or if RSiOH sites do not have a significant effect on lead and phosphate adsorption, then only one type of adsorption site (RFeOH) would need to be considered. The equilibrium constants for the surface acid–base reactions are taken from the published literature for RSiOH sites (Sverjensky and Sahai, 1996) and RFeOH sites (Gunneriusson et al., 1994) (reactions S1–2 and S10–11 in Table 13.4). Surface site densities for RSiOH and RFeOH were set at the value of 2.3 sites/nm2 suggested by Davis and Kent (1990) for oxide minerals. Experimental data were also used in FITEQL to optimize surface site densities, but if optimized site densities were not substantially different from 2.3 sites/nm2, then the value of 2.3 sites/nm2 was used. The ability of equilibrium constants for single sorbate surface complexation reactions from references to simulate our experimental data was evaluated. If the constants from the literature did not provide a good fit, then they were further optimized in FITEQL. This optimization proceeded in two separate steps. First, the equilibrium constants for lead and phosphate surface complexation at
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Table 13.4: Reactions Considered in Aqueous Lead Phosphate System.
P1 P2 P3 P4 S1 S2 S3 S4 S5 S6 S7 S8 S9 S10 S11 S12 S13
Reaction
logKa
Ref.b
Pb5(PO4)3Cl(s)Q5Pb2++3PO3 4 +Cl 2+ 3 Pb5(PO4)3OH(s)Q5Pb +3PO4 +OH + PbHPO4(s)QPb2++PO3 4 +H 2+ – Pb(OH)2(s)QPb +2OH + RFeOH+ 2 2RFeOH+H + RFeOH2RFeO +H RFeOH+Pb2+2RFeOPb++H+ RFeOH+Pb2+2RFeOHPb2+ + RFeOH+PO3 4 +3H 2RFePO4H2+H2O 3 RFeOH+PO4 +2H+2RFePO4H+H2O + 2 RFeOH+PO3 4 +H 2RFePO4 +H2O 2+ 3 + RFeOH+Pb +PO4 +H 2RFePO4Pb+H2O + RFeOH+Pb2++PO432RFeOPbPO 4 +H + + RSiOH2 2RSiOH+H RSiOH2RSiO+H+ RSiOH+Pb2+2RSiOPb++H+ + 2– RSiOH+PO3 4 +H 2RFePO4 +H2O
84.43 76.8 23.81 19.85 7.47 9.51 0.92 9.47 31.13 25.95 18.45
1 1 1 1 2 2 6 6 3 6 6
RFeOH site density ¼ 2.3 sites/nm2 RSiOH site density ¼ 2.3 sites/nm2 Capacitance ¼ 1.28 F/m2
c c
2.4 8.4 1.02
4 4 6
c
5 5 2
a
logK values for surface reactions are for the intrinsic equilibrium constants (Kint). 1: Schecher and McAvoy (1998); 2: Gunneriusson et al. (1994); 3: Nilsson et al. (1992); 4: Sverjensky and Sahai (1996); 5: Davis and Kent (1990); 6: This study. c Reactions were not needed for this study. b
RSiOH sites (reactions S12 and S13 in Table 13.4) were optimized using the datasets for lead and phosphate adsorption to uncoated sand. These values were then held constant as the equilibrium constants for complexation to RFeOH sites (reactions S3–S7 in Table 13.4) were optimized with the datasets for adsorption to goethite-coated sand. Finally, the site (RFeOH and RSiOH) densities and equilibrium constants from single sorbate experiments were combined to assess whether or not they effectively simulate the dual sorbate experimental data. If the combination of single adsorbate reactions did not sufficiently predict dual adsorbate adsorption, then ternary surface complexation reactions (reactions S8 or S9 in Table 13.4) could be added to the model and their equilibrium constants optimized in FITEQL. Precipitation reactions of cholorpyromorphite and lead hydroxide (reactions P1–P4 in Table 13.4) were also included to evaluate their impact on the simulation of lead uptake.
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The software program MINEQL+ (Schecher and McAvoy, 1998) was used to calculate equilibrium speciation based on fixed reactions, constants, and total concentrations. Equilibrium constants for all aqueous phase reactions were those in the MINEQL+ database (Schecher and McAvoy, 1998).
13.3. Results and Discussion 13.3.1. Single Sorbate Adsorption 13.3.1.1. Effect of Goethite Coating on Adsorption The amount of adsorbed lead was always higher on the goethite-coated sand than on the uncoated sand (Fig. 13.2a). However, the iron oxide content was apparently not high enough for goethite to completely dominate adsorption of lead. As seen in SEM images, the sand surface was not completely covered by goethite particles. The quartz surface has sufficient adsorption capacity and affinity for lead that it must be considered in subsequent modeling. The amount of adsorbed phosphate was always higher on the goethitecoated sand than on the uncoated sand (Fig. 13.2b), which indicated that the goethite coating dominated phosphate adsorption. Phosphate adsorption to quartz was so insignificant that adsorption to the quartz was neglected in modeling. Both lead and phosphate adsorb to the goethite-coated sand with patterns typical of the Langmuir isotherm. Because of more favorable adsorption of phosphate than of lead at pH 6, the initial slope of the isotherm is steeper for phosphate than it is for lead.
13.3.1.2. Effect of pH on Lead Adsorption The adsorption of lead on goethite-coated sand was highly pH dependent (Fig. 13.3), increasing from minimal adsorption at pH 4 to nearly complete adsorption at pH 7. Similar pH edges have been observed for lead adsorption to pure goethite (Hayes and Leckie, 1986; Gunneriusson et al., 1994) and are typical of metal cation adsorption to iron oxides (Benjamin, 2002). The SCM, which was used to fit the experimental data of lead adsorption to goethite-coated sand, included two types of adsorption sites, quartz surface hydroxyl (RSiOH) and goethite surface hydroxyl (RFeOH) groups, which were each described by two surface acid–base reactions (S1, S2, S10,
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Figure 13.2: Adsorption of (a) Lead and (b) Phosphate at pH 6, Solid:Liquid Ratio ¼ 50:1 g/l, and 24 h Contact Time. The Ionic Strength was Fixed by NaNO3 at 103 M and 103 M MES Buffer for Lead Adsorption and 102 M and 102 M MES for Phosphate Adsorption.
and S11 in Table 13.4). In the constant capacitance model, a capacitance of 1.28 F/m2 was used based on previous studies (Gunneriusson et al., 1994). The equilibrium constant for lead adsorption to RSiOH groups, with the site density fixed at 2.3 sites/nm2, was optimized by FITEQL as 101.02 using the dataset for lead adsorption to uncoated sand. The quality of the model fit is characterized by a weighted sum of squares divided by the degrees of freedom (WSOS/DF) value of 0.95. The constant for reaction S12 in Table 13.4 was then held constant in modeling lead adsorption to goethitecoated sand. A site density of 2.3 sites/nm2 was also used for RFeOH sites.
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Figure 13.3: Adsorption Edges (a) and Surface Speciation (b) for Lead Adsorption to Goethite-Coated Sand as a Function of pH. Data are Shown with Points and Surface Complexation Modeling Simulations are Shown by the Solid Line. Reaction Conditions: I ¼ 1.2–1.8 103 M NaNO3, Contact Time ¼ 24 h, Solid:Liquid Ratio ¼ 50:1 g/l.
The optimal equilibrium constants for lead adsorption on goethite were determined to be 100.92 and 109.47 for reactions S3 and S4, respectively, with a WSOS/DF value of 2.7. The equilibrium constants are similar to those determined by Gunneriusson and coworkers (1994) for lead adsorption to pure goethite (100.17 and 108.20). Initial modeling was attempted using a single adsorption reaction to goethite, but the model with one reaction could not provide a good fit to the experimental data over the full range of total lead loading. Inclusion of a second reaction significantly improved the quality of the fit over a broader range of total concentration, although the model
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still overpredicts adsorption for the highest lead concentration. The optimization sought the best fit for 5 107 M total lead, because this concentration was also studied in the presence of phosphate. The total site density of the goethite-coated quartz sand was set at 2.3 sites/nm2 for both goethite and quartz surfaces. This value was recommended by Davis and Kent (1990) for mineral oxide surfaces as one for which the results of multiple studies can be compared. When the goethite site density for phosphate adsorption to goethite-coated sand was simultaneously optimized with the equilibrium constant for reaction S7 in Table 13.4, a value of 2.2 sites/nm2 was determined (see next section), and so the value of 2.3 sites/nm2 was used. The site density for goethite is within the range of values used in previous work on lead adsorption to pure goethite of 1.7 sites/nm2 (Gunneriusson et al., 1994) and 7 sites/nm2 (Hayes and Leckie, 1986). The site density of 2.3 sites/nm2 used for quartz is lower than values of 5.9 sites/nm2 (Yee and Fein, 2003) and 5 sites/nm2 (Chen et al., 2006) used in previous studies. Without an independent measurement of site density, the determinations of the site density of surface groups and the equilibrium binding constants are not entirely independent in model derivation based on adsorption data sets. The use of a consistent site density is beneficial for comparison of equilibrium constants from different studies and is important for the development of a self-consistent thermodynamic database (Davis and Kent, 1990). Figure 13.3b shows the simulated surface speciation of lead adsorption to goethite-coated sand. The binding of lead to goethite is stronger than that to quartz and RFeOHPb2+ is the dominant surface species over the experimental pH range. At higher pH (>7.5), RFeOHPb2+ is surpassed by RFeOPb+ as the dominant species. The species RSiOPb+ is never dominant in this model, although its contribution increases with increasing pH.
13.3.1.3. Effect of pH on Phosphate Adsorption The adsorption of phosphate on goethite-coated sand was pH dependent (Fig. 13.4). The phosphate adsorption density decreased with increasing pH. These data were similar to previous observations of phosphate adsorption (Cheng et al., 2004). When the phosphate concentration was low ([P]tot ¼ 5 106 M), complete adsorption on goethite-coated sand was approached when the pH was less than 7. For a total phosphate concentration of 5 105 M, adsorbed phosphate was never more than 35% of the total. Adsorption of phosphate to uncoated quartz sand was insignificant compared with that to goethite-coated sand, therefore, the model used for
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Figure 13.4: Adsorption Edges (a) and Surface Speciation (b) for Phosphate on Goethite-Coated Sand as a Function of pH. Data are Shown with Points and Surface Complexation Modeling Simulations are Shown by the Solid Line. Reaction Conditions: I ¼ 1.2–1.8 102 M NaNO3, Contact Time ¼ 24 h, Ptot ¼ 107 to 104 M, Solid:Liquid Ratio ¼ 50:1 g/l. phosphate adsorption to goethite-coated sand only considered adsorption to sites on goethite (RFeOH). The equilibrium constants of acid–base goethite surface reactions, the site density, and the capacitance were the same as those in the SCM for lead adsorption. Three reactions (S5, S6 and S7 in
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Table 13.4) were included for phosphate adsorption, consistent with previous studies (Nilsson et al., 1992; Gao and Mucci, 2003). Initially the equilibrium constants for reactions S5 and S6 in Table 13.4 were set at values from a previous study of phosphate adsorption to goethite (Nilsson et al., 1992), and the constant for reaction S7 in Table 13.4 was determined by finding the value that provided the optimal agreement between the experimental data and the model simulation (Table 13.4). The optimized logK for reaction S7 in Table 13.4 was 18.45 and the goodness of fit was characterized by a WSOS/DF of 6.5. Because adsorption was already at the maximum for a given total phosphate concentration below pH 7, the quality of the fit was insensitive to the constants for reactions S5 and S6 in Table 13.4; however, optimization of the effect of phosphate adsorption on surface charge for dual sorbate experiments (discussed in the section ‘‘Co-Sorption of Lead and Phosphate in Combined Systems’’) required refinement of the constant for reaction S6 in Table 13.4 to a logK of 25.95. The values of 25.95 and 18.45 for reactions S6 and S7 in Table 13.4, respectively, are lower than the values of 26.38 and 20.61 used by Nilsson and coworkers. is The species RFePO4H is dominant at lower pH and RFePO2 4 increase dominant at higher pH (Fig. 13.4b). Concentrations of RFePO2 4 and those of RFePO4H decrease with increasing pH. The species RFePO4H2 is not significant over the experimental pH range. At higher is more total phosphate concentrations, the formation of RFePO2 4 significant (modeling data not shown). 13.3.2. Co-Sorption of Lead and Phosphate in Combined Systems When lead and phosphate were combined in the same systems, the presence of phosphate increased lead removal for all experiments. The effect was greatest at low pH. At pH 4, lead uptake increased from 6% to 40–54% when the initial conditions of the system were undersaturated with respect to pyromorphite and to 60–86% in systems with initial supersaturation. At higher pH (6 or 7), the presence of phosphate did not have as significant of an impact since lead removal was already approaching 100% in the absence of phosphate. Phosphate sorption in the dual sorbate experiments was measured and was generally consistent with the SCM. Three scenarios were considered as models for the enhancement of lead uptake by phosphate: (1) direct combination of binary surface complexation reactions; (2) combination of binary surface complexation reactions with an additional ternary surface complexation reaction (Table 13.4, reaction S8 or S9); (3) combination of binary surface complexation reactions with potential
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Figure 13.5: Lead Sorption on Goethite-Coated Sand as a Function of pH and Solution Saturation with Respect to Pyromorphite. Fixed Reaction Conditions: I ¼ 1.2–1.8 102 M NaNO3, Contact Time ¼ 24 h, Solid:Liquid Ratio ¼ 50:1 g/l, Pbtot ¼ 5 107 M. Simulations Involve (a) Direct Combination of Binary Adsorption Reactions and (b) Binary Adsorption and Precipitation Reactions. precipitation of pyromorphite and lead hydroxide (Table 13.4, reactions P1 and P4). Figure 13.5a shows the simulation from directly combining surface complexation reactions from lead-only and phosphate-only systems. The model does predict that phosphate addition increases lead adsorption. The increase is caused by the lowering of the surface charge and surface potential (c) from the formation of RFePO4H and RFePO2 4 species, which consequently
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Figure 13.6: Effect of Total Phosphate Loading on Goethite-Coated Sand Surface Charge at pH 4. Simulation Conditions: I ¼ 1.2–1.8 102 M NaNO3, Solid:Liquid Ratio ¼ 50:1 g/L, Pbtot ¼ 5 107 M. Total Phosphate Concentrations Varied from 109 to 103 M (Logarithmic Scale). increases lead adsorption as a result of a more favorable electrostatic component for adsorption (Fig. 13.6). The apparent equilibrium constant for reaction S3 in Table 13.4 (Eq. (13.10)) increases with decreasing surface potential (c) at pH 4 because the intrinsic constant remains fixed (Fig. 13.6). K app ¼ K int eF c=RT ¼
f FeOPbþ gfHþ g fFeOHgfPb2þ g
(13.10)
These simulations (Fig. 13.5a) show a good fit to experimental data at pH 4 and 5. Successfully fitting the lead sorption results at pH 4 did require model refinements beyond those initially attempted. Although only optimization of the constant for reaction S7 in Table 13.4 was necessary to model phosphate adsorption, the resulting effect of phosphate adsorption on surface charge overestimated the increase in lead sorption in the presence of phosphate. Optimization of both reactions S6 and S7 in Table 13.4 provided an equally good fit of phosphate sorption but then provided an improved fit for lead sorption in the dual sorbate experiments. The model predicts decreasing lead adsorption at pH 4.5–5 because the total phosphate concentrations at pH 4.5–5 are less than at pH 4 in order to maintain the same initial saturation indices with respect to pyromorphite (Table 13.3), although this effect remains to be tested experimentally.
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The inclusion of reactions for pyromorphite and lead hydroxide precipitation (reactions P1 and P4 in Table 13.4) only predict higher uptake at the initial saturation index of +1 at pH less than 3.5 and at the initial SI of +5 at pH o 4.5. Pyromorphite was the solid predicted to form at these conditions. At lower initial saturation conditions and higher pH, adsorption was the dominant mechanism for lead uptake. Adsorption reduced dissolved lead and phosphate to concentrations that would no longer be supersaturated. The model simulations are based on equilibrium conditions, but in the actual experimental systems the sorption mechanism may be determined by the relative rates of adsorption and precipitation. The assumption of equilibrium also suggests that there would not be an effect of the order of sorbate addition in dual sorbate experiments; however, this is an effect that should be examined in future work. Inclusion of a ternary surface complexation reaction as a ligand-bridging surface complex (reaction S8 in Table 13.4) did not improve the model fit to the experimental data, nor were other reaction stoichiometries (e.g., metalbridging complexes) helpful for improving the fit. No values for a ternary surface complex improved the model fit. The combination of single-sorbate SCMs already fit the data or slightly overestimated lead uptake, and the inclusion of ternary surface complexes would only increase the predicted uptake. Previous studies of electrophoretic mobility and X-ray absorption spectroscopy show that adsorbed phosphate acts as a reactant to form lead phosphate surface phases that may be highly dispersed on goethite (Weesner and Bleam, 1998). In previous experiments by Ler and Stanforth (2003), the reaction stoichiometries of lead adsorption on phosphated goethite suggested that pyromorphite (Pb5(PO4)3Cl) may precipitate on the surface at conditions that are undersaturated with respect to pyromorphite solubility. The surface precipitation of lead phosphate on the goethite surface is most likely by formation of a new surface phase with the structure of pyromorphite. Analogous research on the co-sorption of zinc and arsenate, a good analog for phosphate, on goethite found that the presence of zinc increased arsenate adsorption and vice versa. The effect was most pronounced at high surface loadings and EXAFS data were used to determine that arsenate and zinc were present in zinc arsenate surface precipitates even at conditions that were undersaturated with respect to zinc arsenate solid phases (Gra¨fe et al., 2004; Gra¨fe and Sparks, 2005). Although these experiments showed no evidence of pyromorphite by SEM and XRD, precipitation of lead phosphate minerals cannot be ruled out. Since the maximum loading of lead on the solid is only 2 mg/kg, the formation of lead phosphate minerals may not be detectable by these
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techniques. The low lead loading would even be challenging to detect by element-specific spectroscopic techniques such as EXAFS. Further investigation with goethite-coated sand on the nature of surface precipitation can focus on conditions of higher initial saturation index and lower pH. The low mass loading of lead on the solid is the result of working with goethite-coated sand, which was selected because of its potential utility as a model geomedium, and valuable information on mechanisms that will apply to this material can be gained from the numerous spectroscopic studies previously conducted with pure goethite.
13.4. Summary Goethite-coated sand can serve as a model porous medium to study the fate of lead in natural soil systems. Lead adsorption to goethite-coated sand can be simulated by a SCM that accounts for adsorption to goethite and quartz surface sites, and phosphate can be modeled using just goethite surface sites. The site densities, reaction stoichiometries, and equilibrium constants for goethite-coated sand are similar to those of pure goethite. The chemical equilibrium model developed in this study can be incorporated into reactive transport models for lead and phosphate transport through sand columns and ultimately to transport in soil and groundwater systems. Over a broad range of solution compositions, phosphate addition enhanced lead removal. Lead sorption can be predicted by the combination of lead and phosphate surface complexation reactions. Ternary surface complexation is not necessary for uptake simulation. The SCM suggests that the dominant sorption mechanism is adsorption, even when the system is initially supersaturated with respect to the lead phosphate mineral pyromorphite. Precipitation mechanisms will become significant at lower pH and higher supersaturation conditions. Since the loading of lead for this study is lower than the detection limits of most techniques for solid phase investigation, the possibility of pyromorphite formation cannot be ruled out based only on macroscopic sorption data.
ACKNOWLEDGMENTS Partial funding for this study was provided by the National Science Foundation (BES #0546219). Liyun Xie has been supported by fellowships through the Washington University School of Engineering and Applied Science, the Cecil Lue-Hing Scholarship, and the Lilia Abron Scholarship.
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The authors acknowledge Claire Farnsworth’s contributions to lead adsorption experiments and Yu Wang’s to phosphate adsorption experiments. The insightful comments of Doug Kent and three anonymous reviewers assisted in the revision of this chapter.
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