Ind. Eng. Chem. Res. 2001, 40, 1615-1623
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MATERIALS AND INTERFACES Adsorption of Cesium, Strontium, and Cobalt Ions on Magnetite and a Magnetite-Silica Composite Armin D. Ebner,† James A. Ritter,*,† and James D. Navratil‡ Department of Chemical Engineering, Swearingen Engineering Center, University of South Carolina, Columbia, South Carolina 29208, and Department of Environmental Engineering and Science, Clemson University, Anderson, South Carolina 29625
Constant pH adsorption isotherms for nonradioactive Cs+, Sr2+, and Co2+ on pure magnetite and a 80% (w/w) magnetite-silica composite were measured at 25 °C over a wide range of metal ion concentrations. The adsorption studies were carried out at four different pH’s: 6, 7, 8, and 9 for Cs+ and Sr2+ and 5, 6, 7, and 8 for Co2+. All of the constant pH isotherms exhibited type I behavior with a saturation capacity that was pH-dependent and increased with increasing pH. The corresponding distribution coefficients increased with increasing pH but decreased with increasing metal ion concentration; they were also 10-1000 times lower than those reported in the literature for more selective but more expensive adsorbents. These two magnetite-based adsorbents also exhibited moderate regeneration conditions, with nearly 90-100% regeneration achieved in most cases at pH values between 1 and 3. A Langmuir model with pH-dependent parameters was also fitted successfully to all of the constant pH adsorption isotherms. This experimental data and the corresponding pH-dependent Langmuir correlation should find considerable use in the design and development of inexpensive fixed-bed adsorption processes for the removal of the radioactive isotopes of Cs+, Sr2+, and Co2+ from aqueous solutions that are produced in nuclear facilities. Magnetite, when encased in silica and placed in a packed column, can also be used as the charging element in high gradient magnetic separation, thereby removing not only metal ions via surface complexation (adsorption) but also nanoparticles of a paramagnetic nature. Introduction As a result of the end of the cold war, the Department of Energy (DOE) began an important program of environmental restoration and decommissioning within the sites that were once used for the production of basic materials for nuclear weapons. Its main objective is to search for alternative solutions for the fast remediation and treatment of all of the radioactive, hazardous, and mixed wastes from basins, burial grounds, landfills, pits, piles, and tanks that resulted as a consequence of almost 40 years of operation. Several intensely radioactive fission products and byproducts are found within these places, among which 90Sr, 137Cs, and 60Co are the most important. 90Sr and 137Cs, for example, are the main fission products of spent fuels, and because of this, they are found in almost all of the radioactively contaminated places, preferentially in the aqueous phases. 60Co, in contrast, does not come directly from the fallout of nuclear fuels but as a consequence of the presence of impurities in the stainless steel that conform nuclear reactors. Much of DOE’s efforts have been focused on removing these highly soluble but dilute species from aqueous * To whom correspondence should be addressed. Phone: (803) 777-3590. Fax: (803) 777-8265. E-mail:
[email protected]. † University of South Carolina. ‡ Clemson University.
waste streams by fixing them into various solid waste forms that can be disposed of in a repository. In this way, the high-volume aqueous streams are transformed from a high-level radioactive waste into a low-level radioactive waste that is much cheaper to treat, e.g., by mixing with cement and storing in specially fabricated casts. However, the removal of these species may not only serve environmental initiatives, but it may also serve as a means of producing useful materials for use in science and industry. It is known, for example, that 137Cs and 60Co are excellent γ sources for medical applications such as instrument disinfection and radiotherapy.1-4 Similarly, 137Cs and 60Co have also proven to be useful as sources for sterilization in the food industry,5-7 and 90Sr, as a source of β radiation, has found applications in medicine, specifically in medical treatments for optical diseases, superficial treatments, and immunotherapy.2,8 Many different processes have been investigated for removing these species from aqueous solutions.9-38 They vary from traditional methods such as precipitation, extraction, ion exchange, and adsorption to relatively less conventional or new ones such as bioaccumulation23-25 and electric-field-assisted techniques.32 Among the inorganic ion-exchange resins and adsorbents investigated for removing these metals, natural and syntheticzeolites,9,10 titanites,11,39,40 silica,40 silicotitanites,12-14 andsilica-orchabazite-supportedmetalhexacyanoferrates15-19
10.1021/ie000695c CCC: $20.00 © 2001 American Chemical Society Published on Web 03/07/2001
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are the most common. However, other materials such as activated carbon,20 coal fly ash,21 bauxite wastes,22 coal ash, and even magnetite41,46,47 have been studied. Magnetite, although not as selective as other adsorbents, is relatively inexpensive and readily available; it is also ferromagnetic, which makes it quite unique as an adsorbent. For example, when fixed within a magnetic field, this magnetic characteristic of magnetite markedly improves its decontaminating capability by allowing the retention of paramagnetic contaminants that are in the form of particles or precipitates. In this regard, investigators48-53 have shown that magnetite can be fixed in resins or inorganic sol-gels and then utilized in a packed column, which not only reduces operational costs but also improves performance. Alternatively, if used only as a traditional metal ion adsorbent, its ferromagnetic property can be used to recover it from a batch adsorption process by high gradient magnetic separation (HGMS). For similar reasons, magnetite has been proposed for the removal of undesirable ferritic compounds from high-level radioactive waste sludge to facilitate vitrification.54 However, a paucity of data exists in the literature that can be used to design fixed-bed adsorption processes utilizing magnetite for the removal of metal ions from aqueous solutions, especially Sr2+, Cs+, and Co2+ ions. Typically, the data in the literature are presented in terms of distribution coefficients that, for the most part, are useful only for batch or fixed-bed adsorption processes designed specifically for treating extremely dilute streams. In reality, the adsorption of metal ions on adsorbents such as magnetite, even from very dilute solutions, exhibit type I adsorption isotherm behavior and thus have a saturation capacity that depends not only on the metal ion concentration but also on the pH of the solution. Therefore, the objective of this paper is to present the adsorption of nonradioactive Sr2+, Cs+, and Co2+ on magnetite and a magnetite-silica [80-20% (w/w)] composite as functions of both the metal ion concentration in solution and the solution pH at 25 °C. The experimental results are correlated with a semiempirical model, based on the Langmuir adsorption isotherm and an empirical dependence of its two parameters on pH. The experimental results are also compared with those of other adsorbents in the literature, in terms of distribution coefficients. The advantages and disadvantages of the use of magnetite and a magnetite-silica composite as effective adsorbents for Sr2+, Cs+, and Co2+ are discussed. Experimental Section Adsorbent Preparation. Two different adsorbents were utilized in the experiments: pure magnetite and an 80% (w/w) magnetite-silica composite.50 Prior to use, sufficient magnetite (Alfa Aesar, 97%) was rinsed with deionized distilled water (pH ) 2, adjusted with HCl), vacuum-filtered in 0.45 mm Tuffryn membrane filters (Gelman Sciences), and vacuum-dried at 110 °C. The 80% (w/w) magnetite-silica composite was prepared as follows. First, 34.68 g of tetraethyl orthosilicate (TEOS; Alfa Æsar, 99%) was combined with 70 mL of deionized distilled water (pH ) 2, adjusted with HNO3) in a sealed 250 mL three-necked flask. The solution was stirred for 16 h at room temperature using a mechanical agitator. Then, 40 g of the rinsed magnetite was added to the TEOS solution, and the solution was stirred for 1 h
Figure 1. Schematic of the experimental setup used to measure the constant pH adsorption isotherms at 25 °C: (A) 500 mL flasks; (B) horizontal shaker bath; (C) nitrogen sparger; (D) pH probe; (E) buret for controlling the pH.
more. Next, during stirring, ammonium hydroxide (2 N) was added dropwise using a dropping funnel and a pH meter until a pH of 8 was obtained; shortly after, the TEOS gelled and effectively trapped all of the magnetite particles contained in the suspension. The gel was removed from the flask, washed, and vacuum-filtered three times with deionized distilled water and then once more with dehydrated ethanol to minimize shrinkage and thus collapse of the microstructure during drying. During each washing, the solution was stirred vigorously for several minutes prior to the filtration step. After washing, the gel was dried at 65 °C for 4 h and then at 110 °C for 4 h more, both using a ramping rate of 0.5 °C/min. The dried sample was calcined in nitrogen at 350 °C for 4 h, using a ramping rate of 5 °C/min, after which it was carefully ground and sieved to a 30-60 Tyler mesh size. Constant pH Adsorption Isotherm Measurements. Constant pH adsorption isotherms were obtained at 25 °C in the experimental setup depicted in Figure 1. The experiments were carried out in four 500 mL flasks (A) placed inside a horizontal shaker bath (B). A total of 5 g (dry weight) of either pure magnetite or an 80% (w/w) magnetite-silica composite was placed in a flask along with 500 mL of solution containing NaNO3 (0.01 N) as the background electrolyte. The solutions were continuously sparged with nitrogen (C), and their pH was continuously monitored with a computer (D) using clear epoxy-bodied, gel-filled combination electrodes (Omega). The pH was maintained at a constant value with HNO3 (0.1 N) or NaOH (0.1 N) solutions using a buret (E). The adsorbent particles [i.e., the magnetite or 80% (w/w) magnetite-silica composite] were suspended in the flask by the effect of the nitrogen sparge and the bath movement. For the first point on the isotherm, a precise dose of 1-4 mL of a 1000 ppm (i.e., mg/L) solution containing the metal ion (metal nitrates, Aldrich) under study was added to the flask. Aliquots of HNO3 or NaOH were continually added to the flask until the pH of the solution stayed at the desired value for more than 1 h, to ensure that equilibrium was achieved. After equilibrium, a 10 mL sample was collected from the flask for analysis and a new precise dose of the 1000 ppm solution was added. Six to seven data points were measured in this way for each isotherm, with isotherms obtained at four different
Ind. Eng. Chem. Res., Vol. 40, No. 7, 2001 1617
Figure 2. Constant pH adsorption isotherms at 25 °C for (1) Co2+, (2) Sr2+, and (3) Cs+ on (a) magnetite and (b) an 80% (w/w) magnetitesilica composite, based on the available surface area of magnetite, and at the following pH’s: pH 9, circles; pH 8, squares; pH 7, triangles; pH 6, crosses; pH 5, diamonds. Symbols represent the experimental data, and lines represent correlations of the experimental data with the pH-dependent Langmuir isotherm model (eqs 1, 3, and 4).
pH’s for each metal: 6, 7, 8, and 9 for Sr2+ and Cs+ and 5, 6, 7, and 8 for Co2+. The samples collected were analyzed by flame atomic absorption spectrophotometry (Perkin-Elmer 3300). Desorption Measurements. Desorption experiments were carried out in a batch mode in 50 mL of solution containing 0.01 N NaNO3, around 25 ppm of either Cs+, Sr2+, or Co2+ and 10 g/L of either pure magnetite or an 80% (w/w) magnetite-silica composite. The solutions were contained in plastic flasks immersed in a shaker bath at 25 °C. Four identical solutions were made for each of the six (i.e., 3 × 2) metal-adsorbent combinations. Initially, 0.5 mL of a 0.1 N solution of NaOH was added to all of the solutions (i.e., a total of 24) to raise the pH. Trial experiments determined that this was sufficient to raise the pH to the point where significant adsorption occurred. A total of 3 days was allowed to reach equilibrium. After equilibrium, one bottle from each metal-adsorbent combination was centrifuged, and its pH and metal concentration were analyzed. Then, for the combinations containing magnetite, 0.5 mL (0.1 N), 0.15 mL (1 N), and 0.75 mL (1 N), and for the combinations containing an 80% (w/w) magnetite-silica composite, 0.65 mL (0.1 N), 0.15 mL (1 N), and 0.75 mL (1 N) of an HCl solution were added to each of the remaining three bottles. Again 3 days was allowed to reach equilibrium. After equilibrium, each solution was then centrifuged, and its pH and metal concentration were analyzed. Model Correlation The experimental results were correlated with the two-parameter Langmuir adsorption isotherm:
qeq )
qmaxkHCeq qmax + kHCeq
(1)
where kH is the Henry’s law constant and qeq and qmax are the equilibrium and the saturation loading, respectively. In linearized form, eq 1 is expressed as
Ceq Ceq 1 ) + qeq kH qmax
(2)
Each experimental isotherm was plotted in terms of eq 2 to obtain the pH-dependent parameters, qmax and kH, from the corresponding slopes and intercepts, respectively. The pH dependence of these parameters was assumed to obey the following empirical expressions:
log kH ) akpH + bk
(3)
log qmax ) aqpH + bq
(4)
The pH-independent parameters (i.e., ak, bk, aq, bq) were obtained from the slopes and intercepts of these functions. Note that eq 3 is actually in the form of a massaction law between the metal species and the protons. Results and Discussion Figure 2 shows the experimental equilibrium adsorption isotherms (symbols) obtained at the different pH’s for Co2+, Sr2+, and Cs+ on both the pure magnetite (parts 1.a, 2.a, and 3.a of Figure 2) and the 80% (w/w) magnetite-silica composite (parts 1.b, 2.b, and 3.b of Figure 2). The equilibrium results for both materials are expressed in terms of the metal loading per unit area of total magnetite, qeq (µmol/m2), to investigate the relative contribution of silica on the adsorption of the metal ions under study. All of the constant pH adsorption isotherms exhibited type I adsorption isotherm behavior,55 with a saturation capacity that depended on pH. In fact, larger adsorption capacities were observed with increasingly larger pH’s for all three metal ions. The increase in capacity with pH resulted from two possible mechanisms: (1) an increase in the number of sites available for metal ion adsorption, because more of the adsorbed protons were thermodynamically forced from the surface to the bulk solution with increasing pH (i.e., at lower bulk concentrations of protons), and (2) an increasingly more negatively charged and thus more electrostatically attractive surface to the cations in the bulk solution, as a result of the decrease in the
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Figure 3. Distribution coefficients for (1) Co2+, (2) Sr2+, and (3) Cs+ on (a) magnetite and (b) an 80% (w/w) magnetite-silica composite at the following pH’s: pH 9, circles; pH 8, squares; pH 7, triangles; pH 6, crosses; pH 5, diamonds. Symbols represent the experimental data, and lines represent correlations of the experimental data with the pH-dependent Langmuir isotherm model (eqs 1, 3, and 4).
number of protons on the surface with increasing pH. However, Co2+ exhibited the largest adsorption capacity on pure magnetite, followed by Sr2+ and then Cs+. The major preference of divalent ions over monovalent ones can be explained by an increased electrostatic contribution to the Gibbs free energy of adsorption for the former.57,58 Less intuitive, however, is the observed preference of Co2+ over Sr2+ because they both have the same valence. According to the solvation model proposed by Sverjensky and co-workers,59,60 the larger ion (Sr2+) should have exhibited the higher adsorption capacity, because larger ions are less affected by repulsive solvation effects when approaching the adsorbent/aqueous interface. However, this theory did not seem to apply here, suggesting that another mechanism besides surface complexation may have played a role. Although not conclusive, a much better explanation to the observed preference of Co2+ over Sr2+ is the similarity in nature and size of cobalt with the iron atoms in the magnetite, wherein the cobalt ions may have also intercalated into the spinel structure of the magnetite. This behavior has been observed, for example, in oxides of other transition metals,42 where Co2+ has been shown to be much less reversible and less influenced by pH and background electrolyte changes than Sr2+. Moreover, this explanation seems quite plausible when considering that Sr2+ and Co2+ exhibited similar adsorption capacities on the 80% (w/w) magnetite-silica composite as shown later. The results with magnetite are also in good agreement with the larger adsorption capacities normally observed for transition and divalent metals over alkaline and monovalent metals.40-45 In general, for all three metal ions, when the 80% (w/ w) magnetite-silica composite was used, the adsorption capacities of all three metal ions were greater than those of the pure magnetite, indicating a significant role of the silica. Recall that all of the loadings were calculated in terms of the surface area of pure magnetite; hence, these larger capacities were simply a consequence of the well-known metal ion adsorptive properties of silica and also its mesoporous structure, which was responsible for the 40-fold increase in the surface area of the composite (aBET ) 84 m2/g) over pure magnetite (aBET ) 2 m2/g). It was also interesting that the 80% (w/w)
magnetite-silica composite compared to the pure magnetite exhibited a significant increase in the adsorption capacity of Sr2+ (parts 2.a and 2.b of Figure 2); this increase was not observed for Co2+ because silica does not have the favorable spinel structure that magnetite has toward Co2+, as explained earlier. As a result, Sr2+ and Co2+ exhibited adsorption properties similar to those of the 80% (w/w) magnetite-silica composite. The differences between pure magnetite and the 80% (w/w) magnetite-silica composite were not as marked for Cs+ (parts 3.a and 3.b of Figure 2), and as alluded to above, Co2+ exhibited the fewest differences (parts 1.a and 1.b of Figure 2). In any case, the capacities of these adsorbents were consistent with results presented elsewhere38-41,46,47 on the adsorption of metal cations by magnetite and silica. Figure 3 shows the experimental results in terms of the distribution coefficient, KD, which is defined as
KD ) 1000[qeqaf/Ceq] (cm3/g)
(5)
where Ceq is the equilibrium concentration of the metal ion in solution (µM), a is the surface area of magnetite (2 m2/g; BET), and f is the weight fraction of magnetite in the adsorbent sample (0.8). In all cases, the distribution coefficients increased with decreasing pH but decreased with increasing concentration of the metal ions in solution. The fact that KD was both pH- and concentration-dependent brings up an interesting problem when trying to compare KD’s to each other in the literature. It is clear from eq 5 that KD is independent of the concentration only in the Henry’s law (infinitedilution) region where the adsorption isotherm is necessarily linear. However, it is shown later that KD is still strongly dependent on pH in this region. Therefore, unless it is proven that KD is measured in the Henry’s law region, little quantitative information can be gained from comparing the KD’s of different adsorbents especially at different pH’s. Nevertheless, qualitatively, it is safe to state here that, in general, magnetite exhibited much lower distribution coefficients (i.e., 100-1000 times lower) for Cs+ and Sr2+ than those reported elsewhere with other inorganic adsorbents, such as chabazites, silicotitanates, zeolites, and hexacyanofer-
Ind. Eng. Chem. Res., Vol. 40, No. 7, 2001 1619 Table 1. Percent of the Metal Ion Recovereda during Desorption Equilibrium Measurements as a Function of pH Cs
Sr
Co
sample
% recovered
% recovered
pH
% recovered
pH
1 2 3 4
0.00 54.2 98.2 87.1
10.83 5.62 3.03 1.85
0.0 98.8 100.2 101.9
8.73 3.61 2.86 1.85
1 2 3 4
80% (w/w) Magnetite-Silica Composite 0.0 9.23 0.0 9.07 0.0 77.8 4.08 93.6 4.83 98.8 80.7 2.98 96.9 2.86 101.5 83.8 1.60 90.1 1.86 101.6
7.50 3.91 2.76 1.89
pH
Pure Magnetite 10.94 0.0 5.78 58.9 3.12 93.2 1.86 96.1
a % recovered ) 100 × (mass adsorbed in sample 1 - mass adsorbed in sample 2)/mass adsorbed in sample 1.
rates.9-19 One of the reasons for these lower values was probably due to the low surface area of magnetite. Another reason was the diminished specificity of magnetite compared to those of other adsorbents for Cs+ and Sr2+, especially in the presence of other competing ions such as Na+. Table 1 shows the desorption results as a function of pH for the three metals and the two adsorbents in terms of the percent of the metal ion recovered relative to that adsorbed in the first sample at high pH. For magnetite, over 90% of the Cs+ and Sr2+ were recovered at pH’s slightly larger but close to 3.0, and essentially 100% of the Co2+ was recovered at similar pH’s. For Sr2+ and Co2+, the results were equally impressive for the 80% (w/w) magnetite-silica composite, with over 90% and essentially 100% recovered at around a pH of 3, respectively. In contrast, only slightly more than 80% of the Cs+ was recovered from the 80% (w/w) magnetite-silica composite even at a pH as low as 1.6. This result was consistent with that reported by Dyer et al.14 for Cs+ and Sr2+ on titanosilicates. They showed that some Cs+ adsorption still occurred even at pH values close to zero but that Sr2+ did not adsorb under these conditions. It is possible that, because of the relatively large ionic size of Cs+, the interaction of this ion with the first sheath of water molecules that covers the ion and normally forbids metal ions from making direct contact with surface sites is especially weak.61 If this is true, Cs+ may have been able to reach the inner Helmholtz layer, thereby making direct contact with the surface site and competing with protons (i.e., in the form of hydronium ions) for surface sites, even at low pH. However, further evidence is needed to support this supposition as to why Cs+ adsorption persists at such low pH values. Despite the above, a noticeable improvement in the adsorption performance of Cs+ and Sr2+ in terms of the distribution coefficients was observed when only 20% (w/w) silica was added to the magnetite, with reasonable reversibility and thus cyclability exhibited at relatively moderate conditions. These results contrast sharply with the strongly acidic conditions required for some of the more selective adsorbents in the literature.16,18 Figure 4 displays the experimental adsorption isotherm data (symbols) plotted according to the linearized Langmuir model depicted in eq 2. In most cases, the experimental results conformed quite well with the model (solid lines), as was also noticed by most of the R2 values reported in Table 2 being close to unity. Table 2 also displays the corresponding values of the Lang-
muir parameters (i.e., maximum adsorption capacities and Henry’s law constants). Notice, however, that the parameter values corresponding to the adsorption of Sr2+ by pure magnetite at a pH of 6 are not included in this table because the isotherm showed some concavity in the lower concentration region which translated into a negatively sloped curve in the linearized Langmuir plot (see Figure 4.2.a). Parts 1 and 2 of Figure 5 show these Langmuir parameters plotted as functions of pH, according to the power law relationships defined in eqs 3 and 4. With the exception of qmax for the adsorption of Cs+ and Sr2+ on magnetite, all of the Langmuir parameters increased with increasing pH. A plausible explanation for the negative relationships observed between qmax and pH for Cs+ and Sr2+ on magnetite was most likely due to these species competing for available sites with Na+ (background electrolyte), especially because magnetite has a very limited surface area available for adsorption. It is also believed that this is the reason for the negative slope observed with Sr+ in Figure 4.2.a, at pH ) 5, which may have indicated that at this pH Sr2+ is the weaker competing species between Sr2+ and Na+. The above explanations are also consistent with Cs+ and Sr2+ exhibiting the positive qmax vs pH trend with the 80% (w/w) magnetite-silica composite, because there was much more surface area available for adsorption. In the case of Co2+, the negative qmax vs pH trend was not observed simply because alkaline metals do not normally compete for the same adsorption sites with transition metals.40,41,56,57 The solid lines in parts 1 and 2 of Figure 5 represent the correlation of the data to eqs 3 and 4. The corresponding pH-independent parameters and R2 values are given in Table 3. In general, the power law models described the pH dependence of the Langmuir parameters quite well. These parameters were used to obtain the solid lines in Figures 2 and 3. In all cases, very good agreement was obtained with the experimental data, even for Sr2+ adsorbed by magnetite at a pH of 6. In this case, the pH-independent parameters obtained solely from the other three Sr2+-magnetite adsorption isotherms at the higher pH’s were used to predict both the adsorption isotherm and KD at a pH of 6. Excellent agreement was obtained in both cases. The Henry’s law constants from the Langmuir model were also used to calculate the infinite-dilution distribution coefficients KoD, by replacing qeq/Ceq in eq 5 with KH. The resulting KoD’s are plotted in Figure 5.3. As stated earlier, in the Henry’s law region, KD is independent of the concentration but not pH. In most cases, KoD increased by an order of magnitude or more, with the pH increasing from 5 to 9. Because of this strong pH dependence, again, it was difficult to make a quantitative comparison with values reported in the literature. Nevertheless, even at the highest pH’s, where the KoD’s were the largest, their magnitudes were still quite low compared to the more selective inorganic ion exchangers.9-19 Despite this drawback, however, it would be advantageous to use one of these relatively inexpensive magnetite-based adsorbents in a layered bed or column-in-series arrangement. For example, the magnetite-based adsorbent could be placed closest to the feed end in a layered bed and used to reversibly adsorb and remove most of the metal ion contaminants in the feed. On top of this layer, or in a
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Figure 4. Constant pH adsorption isotherms at 25 °C for (1) Co2+, (2) Sr2+, and (3) Cs+ on (a) magnetite and (b) an 80% (w/w) magnetitesilica composite in terms of the linearized Langmuir model at the following pH’s: pH 9, circles; pH 8, squares; pH 7, triangles; pH 6, crosses; pH 5, diamonds. Symbols represent the experimental data, and lines represent the correlation of eq 2 with the experimental data. Table 2. Langmuir Adsorption Isotherm Parameters and R2 Values Obtained from the Fit of Eq 2 to the Experimental Dataa qmax (µmol/m2) pH 9 8 7 6 5
Co 3.888 2.118 1.891 0.950
Cos 7.540 3.756 2.890 2.222
Sr
Srs
Cs
Css
1.757 2.073 3.920
9.732 4.461 2.867 1.347
0.476 0.450 0.460 0.703
1.922 1.514 1.075 0.811
kH (L/m2) pH 9 8 7 6 5
Co 0.062 0.030 0.020 0.011
Cos 0.074 0.057 0.038 0.020
Sr
Srs
Cs
Css
0.018 0.010 0.006
0.086 0.038 0.028 0.023
0.026 0.014 0.007 0.003
0.044 0.028 0.018 0.014
R2 pH 9 8 7 6 5
Co 0.969 0.990 0.927 0.926
Cos 0.974 0.989 0.984 0.913
Sr
Srs
Cs
Css
0.939 0.956 0.870
0.866 0.908 0.936 0.993
0.985 0.954 0.913 0.813
0.925 0.991 0.975 0.975
a Subscript s stands for the results corresponding to the 80% (w/w) magnetite-silica composite.
second column in series, a much more selective ion exchanger could be used to irreversibly adsorb and decrease the metal ion concentrations down to acceptable levels. This layered bed or column-in-series arrangement would allow for many moderate regeneration cycles to be carried out with the magnetite before the more selective ion exchanger would require regeneration or replacement. Depending on the application, a significant cost savings could be realized compared to using the more selective ion exchanger alone. Also, recall that a supported magnetite adsorbent, when placed in a magnetic field, could be used to reversibly remove paramagnetic nanoparticles if present in solution.48-53
Conclusions Constant pH adsorption isotherms at 25 °C, for three of the most important radioactive contaminants generated in nuclear reactors and during the production of nuclear weapons, i.e., Cs+, Sr2+, and Co2+, were measured using nonradioactive isotopes on magnetite and an 80% (w/w) magnetite-silica composite at different pH’s ranging from 5 to 9. All of the metal ions exhibited a positive dependence of the adsorption capacity with pH. As a transition metal, Co2+ had the highest adsorption capacity on both the pure magnetite and the 80% (w/w) magnetite-silica composite, followed by divalent Sr2+ and then monovalent Cs+. However, the differences in the adsorption capacities between these metals were not as great for the 80% (w/w) magnetite-silica composite, indicating some sort of specificity of silica in the adsorption of alkaline metals. The distribution coefficients (KD) for these metals were also calculated from the adsorption isotherm data. In all cases, they increased with increasing pH and decreased with increasing equilibrium concentration of the metal ion in solution. At infinite dilution, the KD’s were strongly pH-dependent and ranged between 7 and 163; the values were also considerably higher for the 80% (w/w) magnetite-silica composite. These strong dependencies made it difficult to quantitatively compare these KD’s with those reported elsewhere for these same metal ions adsorbed on different adsorbents; nevertheless, qualitatively, they were 10-1000 times lower, depending on the pH and solution concentration. A positive aspect of these materials, however, was that they exhibited more favorable regeneration conditions compared to the more selective adsorbents, with nearly complete regeneration achieved at pH values between 1 and 3. The Langmuir adsorption isotherm was also fitted to all of the experimental data, with both of the Langmuir parameters taking on empirical pH dependencies. In all cases, a log-linear dependence of the parameters with pH was observed with a remarkably good fit to the data. The corresponding pH-independent parameters were used to backpredict both the adsorption isotherms and the distribution coefficients as functions of pH. The
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Figure 5. Log-linear pH dependence of the Langmuir parameters, (1) qmax and (2) KH, and infinite-dilution distribution coefficients, (3) KoD, for the adsorption of (a) Co2+, (b) Sr2+, and (c) Cs+ on magnetite (open circles) and an 80% (w/w) magnetite-silica composite (closed circles). Table 3. pH-Independent Parameters Obtained from Fitting Eqs 3 and 4 to the Variation in the Langmuir Parameters with pHa Co
Cos
Sr
Srs
Cs
Css
ak 0.2481 0.1919 0.2401 0.1839 0.2991 0.1639 bk -3.2140 -2.6250 -3.9127 -2.7964 -4.2688 -2.8915 2 R 0.9878 0.9628 1.0000 0.8895 0.9990 0.9884 aq 0.1885 0.1706 -0.1742 0.2768 -0.0517 0.1274 bq -0.9327 -0.5440 1.7787 -1.5198 0.0978 -0.8543 R2 0.9413 0.9317 0.8967 0.9896 0.5299 0.9958 a
Subscript s stands for the results corresponding to the 80% (w/w) magnetite-silica composite.
backpredictions were also quite good. In two cases, the pH-dependent Langmuir correlation also revealed the effect of the background electrolyte through the negative relationships observed between qmax and pH with Cs+ and Sr2+. Apparently, these species competed for the same adsorption sites with Na+, the cation of the electrolyte. Overall, the experimental results and corresponding pH-dependent Langmuir correlations reported in this work should find use in designing and developing inexpensive adsorption processes for the removal of Cs+, Sr2+, and Co2+ ions from solution over a wide range of pH’s and metal ion concentrations. Moreover, when magnetite is supported, e.g., as is done here with mesoporous silica, it could be used in a fixed-bed adsorption process; in contrast, operating in a fixed-bed mode with pure magnetite (unsupported) would be nearly impossible because of its small particle size (0.1 µm) and hence high pressure drop. The use of supported magnetite in this sense is of paramount importance
because operating in a fixed-bed mode as opposed to a batch mode, which is the way magnetite is usually used for treating wastewater, always results in a higher separation factor. The magnetic properties of supported magnetite when placed in a column could also be utilized as a charging element in HGMS; in this way, it would remove not only ions but also particles of a paramagnetic nature. Acknowledgment The authors gratefully acknowledge financial support from the National Science Foundation under Grant No. CTS 9985489, and from the Idaho National Engineering and Environmental Laboratory under Contract No. K98564555. Notation a ) BET surface area, m2 g-1 ak ) pH-independent parameter in eq 3 aq ) pH-independent parameter in eq 4 bk ) pH-independent parameter in eq 3 bq ) pH-independent parameter in eq 4 Ceq ) equilibrium concentration in the bulk liquid phase f ) weight fraction of magnetite in the adsorbent kD ) distribution coefficient KoD ) distribution coefficient at infinite dilution kH ) Henry’s law constant qmax ) saturation loading of the adsorbent qeq ) equilibrium concentration or loading in the adsorbent
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Received for review July 26, 2000 Revised manuscript received January 26, 2001 Accepted January 30, 2001 IE000695C
