Adsorption of Phosphate on Manganese Dioxide in Seawater

humic acid were found to suppress the adsorption at low pH, but had no effect in the pH range of seawater. The addition of Ca2+ and Mg2+ in 0.7 M NaCl...
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Environ. Sci. Technol. 1996, 30, 536-541

Adsorption of Phosphate on Manganese Dioxide in Seawater WENSHENG YAO AND FRANK J. MILLERO* Rosenstiel School of Marine and Atmospheric Science, Division of Marine and Atmospheric Chemistry, University of Miami, Miami, Florida 33149-1098

The adsorption of phosphate on manganese dioxide (δMnO2) was determined in seawater as a function of pH, temperature, and salinity. The adsorption in a simple electrolyte solution (0.7 M NaCl) was adequately fitted using the triple-layer surface complexation model, assuming that phosphate formed outer-sphere complexes on the surface of δMnO2. SO42- and humic acid were found to suppress the adsorption at low pH, but had no effect in the pH range of seawater. The addition of Ca2+ and Mg2+ in 0.7 M NaCl was found to enhance the adsorption of phosphate on δMnO2 at pH > 4. This increased adsorption might be due to the changes in both the surface charge of δMnO2 and the solution speciation of phosphate in the presence of Ca2+ and Mg2+. The results suggest that manganese oxides can act as important adsorbents of phosphate in natural waters, as well as in surface sediments, when they are present at higher concentrations than other reactive constituents, such as iron and aluminum oxides.

Introduction As one of the main nutrients, phosphate is essential for aquatic plant growth, but its excess supply can lead to eutrophication. The bioavailability of phosphate in the water column is strongly influenced by the interaction with specific mineral components of bottom sediments (1). A great number of studies have been conducted on the adsorption of phosphate on soils (3, 4), clay minerals (57), and iron(III) (hydr)oxides (7-15). Iron(III) (hydr)oxides, hydrous aluminum oxides, and clay minerals are believed to be important adsorbents for phosphate since all of these substances have positive surface charges at slightly acidic to neutral pH and, therefore, high affinities for anions (16). Manganese(IV) oxides have negative surface charges near neutral pH, i.e., pHZPC = 2.3 (17, 18) and are thought to have a limited capacity to adsorb anions. Little work has been done on the adsorption of anions on manganese oxides (15, 19-21). The surface charge on Manganese(IV) oxides could be reversed by exchange of H+ on the surface for metal ions from solution (22, 23). Kawashima et al. (21) have shown that alkaline earth cations and transition metal * Corresponding author e-mail address: [email protected]; telephone: (305) 361-4707; fax: (305) 361-4144.

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ions cause manganese dioxide to strongly adsorb phosphate between pH 6 and 9. The unique distribution of phosphate across the O2/ H2S interface in the Black Sea has been attributed to the redox cycling of Fe and Mn (28). The minimum concentration of phosphate in the suboxic zone is thought to be a result of adsorption on iron and manganese oxides. Our recent kinetic study indicated that phosphate significantly decreases the rate of oxidation of hydrogen sulfide by manganese dioxide due to its interaction with the MnO2 surface (25). To our knowledge, no study of the adsorption of phosphate on manganese dioxide has been made in seawater. In this study, we examine the adsorption of phosphate on manganese dioxide (δMnO2) in simple salt solutions and seawater. The effects of pH, salinity, the major ions of seawater, and humic acid on the adsorption have been investigated.

Experimental Section Materials. The manganese dioxide (δMnO2) was prepared according to the procedures given in Murray (22). The δMnO2 suspension used in the study had been aged for years. The aged δMnO2 had a surface area of 206 m2/g, as determined by the BET method. Goethite was prepared according to Schwertmann and Cornell (26) and had a surface area (BET) ) 48.8 m2/g. All chemicals used were analytical reagent grade and contained negligible phosphate impurity. The pure water used was deionized MilliporeSuper Q (18 ΩW). The phosphate stock solution was prepared from KH2PO4. The seawater used was Gulf Stream surface water collected 10 miles off the coast of Miami and filtered through a 0.45 µm Millipore filter. For the salinity dependence studies, the filtered seawater was diluted with Milli-Q water. Adsorption Measurements. The adsorption of phosphate on δMnO2 was measured in pure water, 0.7 M NaCl, and seawater as a function of pH (2-8.5). Pure water and NaCl solutions contained 2 mM NaHCO3 as a pH buffer. The initial concentrations of phosphate and δMnO2 in most experiments were 10 µM and 1 mM, respectively. The pH of the solutions was adjusted by the addition of HCl or NaOH. The kinetic runs were carried out in a 250 mL waterjacketed glass cell, and the temperature was controlled to 0.05 °C with a Forma temperature bath. The solutions were stirred at the same rate during all kinetic runs. Preliminary kinetic experiments indicated that the initial adsorption reaction was rapid and reached equilibrium (∼95%) after 10 h. Equilibrium experiments were performed under room temperature (23 ( 1 °C) in 250 mL poly(ethylene) bottles. The solutions were allowed to equilibrate in a shaker for 24 h. After equilibration, the pH of the suspension was measured with an Orion glass pH electrode and an Ag/ AgCl double-junction reference electrode using a Metrohm pH meter. The electrodes were calibrated with Tris buffers (27, 28). The samples were then filtered through 0.45 µm filters, and the concentrations of phosphate in the filtrates were determined by colorimetric analysis (29). Modeling of the Adsorption Processes. A number of surface complexation models have been developed over the past 2 decades to describe the equilibrium of adsorp-

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TABLE 1

Surface Parameters of δMnO2 Used in Modelinga surface area surface site density inner-layer capacitance outer-layer capacitance

206 m2/g 18 sites/nm2 2.4 F/m2 0.2 F/m2 log K

SOH2+ T SOH + H+ SOH T SO- + H+ SOH + Na+ T SO--Na+ + H+ a

1.6 -6.2 -3.5

Ref 18.

FIGURE 2. Adsorption isotherms in seawater at pH 8.0, [δMnO2] ) 0.2 mM: (a) Langmuir and (b) Freundlich. FIGURE 1. Uptake kinetics of phosphate on δMnO2 in seawater at pH 8.0 (inserted figure is for the enlargement of initial uptake period). [δMnO2] ) 1 mM, [PO43-]0 ) 10 µM.

tion-desorption reactions at a mineral surface. These models have recently been reviewed by Davis and Kent (30). In this study, we used the triple-layer model to determine equilibrium binding constants for phosphate interactions with manganese dioxide. Detailed discussions and formulations of the triple-layer model are given elsewhere (31-33). The model computations were facilitated by using the computer program HYDRAQL (39). Model-adjustable parameters and the surface equilibria and constants (18) are listed in Table 1.

Results Adsorption Kinetics. The removal of phosphate by δMnO2 in seawater (pH ) 8.0) was followed over a time scale of 50 h (Figure 1). The adsorption reaction was initially very rapid and was 60% complete in 5 min. The rate of phosphate adsorption decreased with prolonged reaction time. The reaction reached near equilibrium (95%) after 10 h (Figure 1). This rapid uptake has implications for understanding the dynamics of phosphate adsorption in natural waters when comparing the time scales of the formation of MnO2 from the oxidation of Mn(II) and the time scales of the physical transport processes. Long-term kinetic results of phosphate adsorption on goethite showed a slow release of phosphate after several days as a result of coagulation and rearrangement processes (11). Within the time scale of our experiments (about 2 days), no oscillatory uptake kinetics was observed (Figure 1). If the initial rate of adsorption was simply regarded as pseudofirst order over a period of 5 min (Figure 1), the temperature dependence of the initial rate constants gave an apparent activation energy of 9 kJ mol-1.

Adsorption Isotherms. Equilibrium experiments of phosphate adsorption were carried out at initial phosphate concentrations between 2 and 80 µM in seawater at pH 8.0. The concentration of δMnO2 in the solutions was 0.2 mM. Adsorption data plotted in both Langmuir and Freundlich equations exhibit three linear segments, instead of a single straight line, over the concentration range examined (Figure 2). Similar results were found for the adsorption of Zn and Cd by δMnO2 (35). These observations are consistent with the existence of a variety of binding sites with a wide spectrum of binding energies (18). At low adsorption densities, the surface sites with highest energy were occupied first. Segments 1, 2, and 3 in the Langmuir isotherm (Figure 2a) gave site concentrations of 2.2, 6.2, and 6.4 µM for 0.2 mM MnO2, corresponding to initial concentrations of phosphate of 2-6, 6-25, and 25-80 µM, respectively. The absorbed PO43- in the following experiments ([MnO2] ) 1 mM, initial [PO43-] ) 10 µM) occupied the high-energy sites (group 1 in Figure 2a). Adsorption of Phosphate in NaCl Solution and Seawater. The adsorption of phosphate on δMnO2 was measured in 0.7 M NaCl solution and seawater as a function of pH. The maximum in the adsorption of phosphate occurred at low pH, and the adsorption decreased with increasing pH (Figure 3). The degree of phosphate uptake by δMnO2 in seawater (S ) 35) was greater at pH > 4 and lower at pH < 4 than that in 0.7 M NaCl (Figure 3) due to the influence of the major ions of seawater. The phosphate adsorption in seawater was found to be independent of salinity in the range of S ) 5-35 (Figure 4). Effect of Ca2+, Mg2+, and SO42-. To examine the effect of the major ions Ca2+, Mg2+, and SO42- in seawater on phosphate adsorption, a series of measurements was made in 0.7 M NaCl with the addition of an individual major ion

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FIGURE 3. Adsorption of phosphate by δMnO2 in 0.7 M NaCl and in seawater. [δMnO2] ) 1 mM, [PO43-]0 ) 10 µM.

FIGURE 6. Effect of Mg2+ on the adsorption of phosphate by δMnO2 in 0.7 M NaCl. [δMnO2] ) 1 mM, [PO43-]0 ) 10 µM.

FIGURE 4. Adsorption of phosphate by δMnO2 in seawater as a function of salinity. [δMnO2] ) 1 mM, [PO43-]0 ) 10 µM.

FIGURE 7. Effect of SO42- on the adsorption of phosphate by δMnO2 in 0.7 M NaCl: solid line, model fit for 0.7 M NaCl; dashed line, model fit for 0.7 M NaCl + 0.029 M Na2SO4 considering the formation of surface complex SOH2+-HSO4-. [δMnO2] ) 1 mM, [PO43-]0 ) 10 µM.

Effect of Humic Acid. Despite its low concentration in seawater, organic matter such as humic acid is known to modify the surface physicochemical properties of particles (42). The effect of humic acid (10 mg/L) on phosphate adsorption on δMnO2 was examined in 0.7 M NaCl. Humic acid was found to suppress phosphate adsorption significantly below pH 5 and showed no effect above pH 6 (Figure 8). The maximum effect occurred in the range of pH 3-4, which coincided with the region pKa ) 3.5-4.0 of humic acid (43). The results suggest that humic acid is unlikely to influence the adsorption of phosphate on δMnO2 considerably in the pH range of most natural waters. FIGURE 5. Effect of Ca2+ on the adsorption of phosphate by δMnO2 in 0.7 M NaCl (dashed line indicates the model fit considering the formation of surface complex SO--Ca2+). [δMnO2] ) 1 mM, [PO43-]0 ) 10 µM.

at its oceanic level. The results are shown in Figures 5-7. The results in NaCl are shown in each figure for comparison. Ca2+ and Mg2+ were found to significantly enhance the adsorption of phosphate on δMnO2 above pH ≈4 and showed little effect at low pH (Figures 5 and 6). The SO42ion was found to suppress phosphate adsorption at low pH and had no effect above pH 4 (Figure 7).

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Discussion The triple-layer model was used to quantitatively describe the phosphate adsorption by MnO2. We started with the experimental data in 0.7 M NaCl solution. In order to model the data, choice of the type of surface complex, stoichiometry, and magnitude of the formation constant for each surface complex is required. The association of anions with oxide surface sites could be described by various reactions, such as the formation of mono- and bidentate, inner- and outer-sphere complexes. The possible reactions between phosphate and the δMnO2 surface (SOH) could be presented

FIGURE 8. Effect of humic acid (10 mg/L) on the adsorption of phosphate by δMnO2 in 0.7 M NaCl. [δMnO2] ) 1 mM, [PO43-]0 ) 10 µM.

as follows:

inner-sphere complexes SOH + PO43- + H+ T SPO42- + H2O

(1)

SOH + PO43- + 2H+ T SHPO4- + H2O

(2)

SOH + PO43- + 3H+ T SH2PO4 + H2O

(3)

2SOH + PO43- + 2H+ T S2PO4- + 2H2O

(4)

outer-sphere complexes SOH + PO43- + H+ T SOH2+-PO43-

FIGURE 9. Adsorption of phosphate in (a) 0.7 M NaCl (2 mM NaHCO3) and (b) H2O (2 mM NaHCO3): solid circles, experimental data; solid line, model fit considering inner-sphere SPO42- species; dashed line, model fit considering a combination of outer-sphere SOH2+H2PO4- and SOH2+-HPO42- species.

(5)

TABLE 2

SOH + PO43- + 2H+ T SOH2+-HPO42-

(6)

Equilibrium Constants for Anion Adsorption on δMnO2 Using the Triple Layer Model

SOH + PO43- + 3H+ T SOH2+-H2PO4-

(7)

2SOH + PO43- + 2H+ T (SOH2+)2-PO43-

(8)

log K

The choice of reactions for the surface species is not necessarily unique. Two or more sets of different model parameters have been shown to provide equally good fits to titration and adsorption data (36, 37). We first fit our adsorption data using the inner-sphere reactions. The solid line shown in Figure 9a is the model fit by assuming the formation of SPO42- surface complex (eq 1). The model adsorption edge was too steep compared to the experimental data. Model calculation including carbonatesurface complexation did not improve the fitting result. Choice of the surface complexes SHPO4-, SH2PO, and S2PO-, (eqs 2, 3, and 4, respectively) resulted in even steeper adsorption edges because the formation of these species involved the uptake of more protons than SPO42- species. The adsorption data in H2O and the model calculation using the surface species SPO42- and the formation constant obtained in 0.7 M NaCl (log K ) 29.0, Table 2) are compared in Figure 9b. The model based on the formation of an inner-sphere complex gave similar results for H2O and 0.7 M NaCl (solid line in Figure 9b), while the experimental data showed that the adsorption of phosphate in H2O was much less than that in 0.7 M NaCl. These results suggest that phosphate is unlikely to form inner-sphere complexes (in terms of the triple-layer model) with the surface sites of δMnO2. Surface spectroscopic work is needed to examine

H+

SOH + PO4 + T SPO4 + H2O SOH + PO43- + 3H+ T SOH2+-H2PO4SOH + PO43- + 2H+ T SOH2+-HPO42SOH + SO42- + 2H+ T SOH2+-HSO43-

2-

29.0 25.1 19.6 5.7

the coordination and structure of adsorbed phosphate on the δMnO2 surface. The adsorption data in 0.7 M NaCl from pH 3 to 8.5 were fit successfully by considering the formation of outer-sphere complexes SOH2+-H2PO4- and SOH2+-HPO42- (dashed line in Figure 9a). The formation constants obtained from the modeling are listed in Table 2. Inclusion of the species SOH2+-PO43- (eq 5) in the model calculations resulted in an increase of phosphate adsorption at pH > 8.0, which was inconsistent with the experimental results. Model simulations for binuclear surface complex (SOH2+)2-PO43(eq 8) did not improve the fit. The model overestimated the amount of phosphate adsorbed on δMnO2 below pH 3.0. One possible explanation is that the number of surface sites was reduced by aggregation processes (12) due to the proximity to the isoelectric point of the solid (pHZPC = 2.3 for δMnO2). Another possibility is the presence of anion impurities and/or Cl- in the NaCl solution that competed for the surface sites with phosphate at low pH. The model did predict a decreasing phosphate adsorption in H2O by use of the formation constants for SOH2+-H2PO4- and SOH2+-HPO42- determined for 0.7 M NaCl, but did not reproduce the experimental data (dashed line in Figure 9b).

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FIGURE 10. Distribution of solution speciation of phosphate in (a) 0.7 M NaCl, (b) 0.7 M NaCl + 0.011 M CaCl2, and (c) 0.7 M NaCl + 0.055 M MgCl2.

The adsorption model was also used to examine the effect of Ca2+, Mg2+, and SO42- in the mixed electrolytes by use of the complexation constants for SOH2+-H2PO4- and SOH2+-HPO42- (Table 2) determined for the NaCl solution. The addition of Ca2+ and Mg2+ caused changes in both the solution speciation of phosphate (Figure 10) and the surface speciation of δMnO2:

SOH + Ca2+ T SO--Ca2+ + H+

(9)

SOH + Mg2+ T SO--Mg2+ + H+

(10)

Ca2+ and Mg2+ competed for the surface sites with phosphate, resulting in both a decrease in the availability of surface sites and a reverse of the surface charge. Suppression of the adsorption of the simple organic acids (anions) on goethite in the presence of Ca2+ and Mg2+ has been partly attributed to the competition between Ca2+ and Mg2+ and the organic acids for surface sites (39). However, Tipping (40) found that Ca2+ and Mg2+ enhanced the adsorption of humic acids on goethite. The increase in the adsorption of arsenate and phosphate by Mn(IV) oxides at near-neutral pH in the presence of divalent cations has been attributed to a reverse of the surface charge from negative to positive by exchange of H+ on the Mn(IV) oxide surface for divalent cations (20, 21). Our results suggest that the change in the surface charge has a greater effect on the adsorption behavior of phosphate than does the reduction of the number of surface sites due to the interactions between Ca2+ and Mg2+ and the δMnO2 surface. However, the model fit by use of the formation of surface complexes SO--Ca2+ and SO--Mg2+ (eq 9 and 10) could not reproduce the experimental data. By using the formation constant log K ) -3.75 for SO--Ca2+, the model

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predicted an increase in adsorption in the pH range 5.57.0, but substantially underestimated the adsorption above pH 7 (dashed line in Figure 5). Use of a higher formation constant K did not improve the overall fit, nor did the inclusion of the species SO--CaOH+, SO--CaHCO3+, SOH2+-CaCO30, SO--CaOH+, and (SO-)2-Ca2+ (41) in the model calculations. The model gave an even worse fit for the adsorption data with the addition of Mg2+ (results not shown). These results indicate that the interaction between Ca2+ and Mg2+ and the δMnO2 surface alone cannot account for the increased adsorption of phosphate. As mentioned earlier, the addition of Ca2+ and Mg2+ to NaCl solutions also changed the phosphate speciation (Figure 10). Ca2+ and Mg2+ form ion pairs with phosphate at pH 4.5 and 4.0 (Figure 10b,c), respectively. At almost the same values of pH, the adsorption of phosphate started to increase in the presence of Ca2+ and Mg2+ (Figures 5 and 6). These results suggest that Ca2+-phosphate and Mg2+phosphate species might be adsorbed on the δMnO2 surface leading to the enhanced adsorption of phosphate. By considering the adsorption of Ca2+-phosphate and Mg2+phosphate species, such as SO--CaH2PO4+, SOH2+CaPO4-, SO--CaHPO40, SO--MgH2PO4+, SOH2+-MgPO4-, SO--MgHPO40, the model did predict an increase in adsorption, but the pH dependence of the increase did not reproduce the experimental data. Hawke et al. (14) encountered similar difficulties in modeling the enhanced adsorption of phosphate on goethite in the presence of Ca2+ and proposed that the increased adsorption was caused by precipitation of a calcium phosphate phase on the particle surface. Surface spectroscopic study is needed to confirm this mechanism. The formation of CaCO3(s) in the NaCl solution is unlikely under our experimental conditions due to the low degree of saturation. The small effect of Ca2+ and Mg2+ at pH < 4 (Figures 5 and 6) was due to the fact that Ca2+ and Mg2+ were not adsorbed on the δMnO2 surface and did not affect the solution speciation of phosphate (Figure 10b,c) at low pH. The combining effect of Ca2+ and Mg2+ in seawater was smaller than expected (Figures 3, 5, and 6), which might be partly due to the interactions between Ca2+, Mg2+, and SO42in seawater. In addition, the interactions of the anions in seawater, such as SO42-, HCO3- (CO32-), B(OH)4-, and F-, and the δMnO2 surface might be enhanced by Ca2+ and Mg2+ (like phosphate) and decreased the phosphate adsorption to some extent. The addition of SO42- in NaCl did not change the solution speciation of phosphate. The suppression of phosphate adsorption at low pH by SO42- apparently was due to the competition between SO42- and phosphate for particular surface sites. The experimental data were fit successfully by use of the formation of the SOH2+-HSO4- complex (dotted line in Figure 7):

SOH + SO42- + 2H+ T SOH2+-HSO4-

(11)

with the formation constant log K ) 5.7. Model calculations showed that SO42- occupied 12-0.2% of surface sites at pH 2.0-4.0 and did not associate with the particle surface above pH 4.0. The suppressing effect of SO42- in seawater was less than that in NaCl solution (Figures 3 and 7) due to the interactions between SO42- and Ca2+ and Mg2+ in seawater and the resultant decrease in the activity of SO42-. In the pH range of most natural waters, the effect of SO42- on adsorption is not expected to be important. Surface

FIGURE 11. Comparison of phosphate adsorption by δMnO2 and goethite in seawater. [δMnO2] ) 1 mM, [r-FeOOH] ) 1 mM, [PO43-]0 ) 10 µM.

complexation models are not applicable to the adsorption of macromolecular polyelectrolytes such as humic acid (44). It is likely that humic acid, as SO42-, decreased the phosphate adsorption by competing for the surface sites. The unsuccessful modeling of the phosphate adsorption in different media (Figure 9b) and the effect of Ca2+ and Mg2+ indicates that we do not fully understand the adsorption processes on MnO2 and/or there are defects in the surface complexation models (as least for MnO2). More studies on the ion interactions with the MnO2 surface are needed to improve our understanding of the adsorption mechanism. The adsorption of phosphate on δMnO2 and goethite in seawater is compared in Figure 11. The results for goethite were similar to a previous study by Hawke et al. (14). The concentrations of δMnO2 and goethite used were 1 mM, which corresponded to 510 and 18 µM surface sites (using a site density of 2.6 sites/nm2 for goethite) (38), respectively. At pH > 7, goethite had a greater affinity for phosphate than did δMnO2 per surface site. However, on the basis of the total concentration of particles, the difference in the degree of phosphate adsorption on δMnO2 and goethite became smaller (∼15% at the pH range of seawater; Figure 11) since δMnO2 has a larger surface area and site density than goethite. Model calculations were carried out to evaluate the relative significance of phosphate adsorption by MnO2 and R-FeOOH. When the concentration of PO43is much higher than the available surface sites of oxides, the importance of MnO2 only depends on its abundance, since the goethite surface has been saturated. The results suggest that manganese oxides can act as important adsorbents of phosphate in natural waters, as well as in surface sediments, when they are present at higher concentrations than other reactive constituents, such as Fe and Al oxides.

Acknowledgments This work was supported by the Office of Naval Research. We thank L. S. Balistrieri and J. W. Murray for helpful discussions on the HYDRAQL program.

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Received for review April 26, 1995. Revised manuscript received September 1, 1995. Accepted September 5, 1995.X ES950290X X

Abstract published in Advance ACS Abstracts, December 1, 1995.

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