Article pubs.acs.org/jchemeduc
Adsorption of Oxy-Anions in the Teaching Laboratory: An Experiment To Study a Fundamental Environmental Engineering Problem Mitch D’Arcy,*,† Florence Bullough,† Chris Moffat,†,‡ Edoardo Borgomeo,† Micheal Teh,† Ramon Vilar,‡ and Dominik J. Weiss†,§ †
Department of Earth Science and Engineering, Imperial College London, London SW7 2AZ, United Kingdom Department of Chemistry, Imperial College London, London SW7 2AZ, United Kingdom § School of Earth Science, Stanford University, Stanford, California 94305, United States ‡
S Supporting Information *
ABSTRACT: Synthesizing and testing bicomposite adsorbents for the removal of environmentally problematic oxy-anions is high on the agenda of research-led universities. Here we present a laboratory module successfully developed at Imperial College London that introduces the advanced undergraduate student in engineering (chemical, civil, earth) and science (chemistry, materials, earth science) to several fundamental principles associated with this research area in a simple, engaging and safe way. This includes (i) the synthesis of inorganic bicomposite sorbents, (ii) the evaluation of the adsorption−removal process, and (iii) the analysis of the sorbate; all underpinned by theory. We devise an experiment using phosphate oxyanions and an iron−titanium oxide bicomposite sorbent, which is simple to synthesize. The adsorption of phosphate solutions of varying concentration is tested and assessed at pH 5 and 9. Phosphate concentrations at equilibrium are analyzed using UV−vis spectroscopy to plot adsorption isotherms and compare the Langmuir and Freundlich models. This topical introduction to environmental engineering is an excellent opportunity to investigate adsorption processes. The complexity of data interpretation can be tailored to a range of abilities, and the topics addressed in the experiment are relevant starting points for further exploration of environmental geochemistry, pollution control, element transport, and adsorption. We present ready-to-use spreadsheets for the students to facilitate data analysis. KEYWORDS: Upper-Division Undergraduate, Analytical Chemistry, Chemical Engineering, Environmental Chemistry, Laboratory Instruction, Physical Chemistry, Hands-On Learning/Manipulatives, Aqueous Solution Chemistry, Geochemistry
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Adsorption is the decontamination method of choice in most cases, and this is a fundamental physicochemical process in environmental science and engineering, geology, and inorganic chemistry, including soil science, pollution control, mineralogy, solid surface chemistry, element transport, minerals processing and refinement, and acid mine drainage. Adsorption is a surface process in which an adsorbate (e.g., an oxy-anion pollutant) binds with receptor sites on the surface of a solid sorbent (e.g., a metal oxide). After a variable reaction time, equilibrium is established, where the sorbent sites are saturated and the adsorbate concentrations on the solid surface and remaining in solution are constant. Binding may be via electrostatic attraction or a range of chemical bonds, forming inner- and outer-sphere complexes.4 Mixed metal oxide bicomposites are of special interest during the development of novel sorbents as they have the capability to perform numerous functions essential for decontamination (e.g., catalysis, oxidation, adsorption) at the same time and are therefore multifunctional materials.5 For example, the TiO2−
he presence of oxy-anions in natural waters can cause a range of environmental problems. For example, high concentrations of phosphate, PO43−, and nitrate, NO3−, disrupt aquatic ecosystems. Known as eutrophication, this problem is caused by the fertilizing properties of the oxy-anions, which normally limit the populations of algae and phytoplankton in rivers and lakes. Excessive concentrations of phosphate and nitrate cause “algal blooms”, which deplete dissolved oxygen and destroy aerobically respiring organisms that cannot tolerate hypoxia, such as fish and shellfish. Algal blooms may also be toxic themselves and promote the growth of toxic bacteria.1 Another harmful oxy-anion is arsenate, AsO43−. Dissolution of arsenate in the groundwater of the Bengal basin has caused the largest mass poisoning in human history, with up to 200 million people drinking poisoned water and suffering various diseases and cancers as a result.2 Other examples of environmentally problematic oxy-anions are uranyl, pertechnetate, and selenate.3 Often these metal species are leached into natural waters from buried waste and the tailings of mining operations, and so cleanup technologies are also relevant for the sustainability of industrial operations. © 2013 American Chemical Society and Division of Chemical Education, Inc.
Published: November 7, 2013 505
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how the sorbent surface behaves. We use UV−vis spectroscopy to measure phosphate concentration and challenge the students to prepare their own sorbent. This enables them to consider the chemical properties that make it a suitable product and exposes them to a wider range of laboratory techniques. We also introduce bicomposite adsorbents, which are becoming increasingly important in adsorption research. Other published experiments, which focus on the adsorption process in itself, lack important steps in data analysis. For example Tribe and Barja introduce an experiment for phosphate adsorption onto an iron oxide mineral.11 They include sorbent synthesis and vary pH; however, they limit data analysis to the Langmuir isotherm and do not handle uncertainty. We feel it is important for students to learn how to handle uncertainty in their data, and we compare the Freundlich and Langmuir isotherm models to encourage the student to think critically about what these models really represent.
Fe2O3 bicomposite mineral sorbent synthesized in this exercise was originally designed to treat drinking water contaminated with arsenic. It oxidizes AsIII to AsV by catalyzing photooxidation6 and quickly uptakes a significant mass of dissolved AsV.7,8 Given the increasing importance of multifunctional materials in research and industry, this mixed oxide is an ideal example for students to synthesize; it is inexpensive, can be prepared quickly using a variety of laboratory techniques, and has a very measurable performance. Although designed for arsenic removal,6 the sorbent can be applied to a range of oxyanions. Phosphate is used in this experiment because it is safe for students to handle and equally relevant in environmental geochemistry. Phosphate can be effectively removed from contaminated water by adsorptive filtration using metal oxides.9 It quickly adsorbs from solution onto mixed Fe2O3−TiO2 oxide particles in solid suspensions, with a measurable decrease in remaining concentration over durations of minutes to tens of minutes as shown in a recent research project conducted in our laboratories.7 Additionally, the concentration of phosphate in solution can be inexpensively and easily measured using UV− vis spectroscopy, therefore teaching the student the very basics of analytical chemistry. In summary, the adsorption of dissolved phosphate onto the bicomposite is a simple, safe, and repeatable experiment, which neatly demonstrates mineral surface reactions and introduces the topic of clean water technologies. The experiment also provides opportunities to introduce the importance of variable control, precision, time and space organization in the lab, repeat measurements, and equipment cleanliness. The use of key lab apparatus such as mechanical pipets, syringe filters, and pH meters can be practiced many times by students during solution preparation and sampling. The experiment is adaptable in difficulty, complexity, sorbents, solutes, and analytical techniques. It can be completed by individuals or groups in as little as a few hours or extended to longer time periods for a more detailed kinetic inspection as we do in our laboratory undergraduate course at Imperial College London. Once students have completed their experiment and carried out measurements, they can graphically determine the type of adsorption reaction with respect to phosphate using traditional models and Microsoft Excel. Example student data taken from one of our lab classes and calculation spreadsheets are provided as Supporting Information if it is not possible to do this using data collected by students. The experiment should be placed in a degree program according to the syllabus; for example, we consider it appropriate for second- and third-year undergraduate students in science (chemistry, materials) and engineering (civil, environmental, chemical) in the United Kingdom and the United States. The experiment presented here forms part of a successful course conducted at the advanced undergraduate level for Earth Science and Engineering students at Imperial College London. A small number of adsorption experiments have been suggested for university courses in the past. Some of these place more emphasis on analytical techniques for measuring adsorbate concentration; for example, Schuttlefield et al. outline an experiment for using ATR-FTIR spectroscopy to investigate sulfate adsorption onto TiO2.10 This may be more appropriate if the course co-ordinator does not want to cover the adsorption process in detail. Our experiment differs by varying the pH of the reaction mixture so that students can understand
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BACKGROUND INFORMATION AND RELEVANT THEORY
Bicomposites and Mineral Synthesis
Bicomposite sorbents are multifunctional materials that are developed to tackle a range of water contamination problems. They can perform oxidative, catalytic, and high-capacity adsorptive functions as outlined above. In addition, their synthesis is inexpensive, relatively quick and easy to automate, repeatable, yet adaptable (e.g., allowing easy control of particle and pore size in the product, or the ratio of constituent minerals and metals in its composition). Solid adsorbents have the advantages of being easy to store, transport and apply to water, safe to handle, and easily removed from solution by filtration. They have a small environmental footprint and are recyclable. Many published sorbent synthetic routes aim to achieve materials with extremely porous surface structure and as a result involve advanced mineralization techniques. Simplifying the preparation process only marginally reduces the adsorptive capacity of the product, meaning both the synthesis and application of a multifunctional sorbent is possible in a taught laboratory class.7 Figure 1 provides a general overview of the adsorbent synthesis. Students prepare the sorbent by precipitating ferric oxide onto a solid titanium dioxide substrate, from a solution of Fe(NO3)3 in ethanol. The volumes and concentrations recommended produce mixed oxide particles in which both mineral phases are exposed at the surface with a Fe/Ti ratio of
Figure 1. (A) Schematic diagram showing the synthesis steps for the Fe2O3−TiO2 adsorbent. (B) Photograph of the dried and crushed adsorbent. 506
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about 1:1.6 An extension to this lab class would be to verify the surface mineralogy using X-ray diffraction or experiment with different iron nitrate concentrations. Precipitating iron oxide from an ethanol solution is quick and safe, and the ethanol can be evaporated easily using a rotary evaporator (or left overnight in an open beaker inside a fumehood). Before evaporation of the alcohol, the mixture is stirred using a magnetic stirrer and sonicated using an ultrasonic water bath; which, coupled with pipetting, dissolving a salt, and the measurement of precise volumes using flasks and masses using a mass balance, exposes the students to a range of laboratory skills. Sonication improves the porosity of the product but can be skipped if the equipment is not available. Similarly, magnetic stirring may be substituted with manual stirring with a spatula if needed. After ethanol evaporation the solid product is crushed and calcined at 300 °C, producing a dry product and allowing students to use a furnace. The synthesis procedure outlined in the Supporting Information should produce a uniform, bright red powder with a particle surface area of approximately 130 m2 g−1. The product can be stored in any dry container without degrading, is nontoxic, and is easy to weigh. Interested students are referred to Zhou et al.6 and D’Arcy et 7 al. for a more detailed explanation of how the metal oxide components in the bicomposite behave with respect to arsenic adsorption. The sorbent is appropriate for phosphate adsorption as well,7 but arsenic uptake illustrates the multifunctional benefits of using a mixed oxide, which is of particular interest for engineering applications of adsorption. To summarize, the titanium phase catalyzes the photo-oxidation of AsIII to AsV (the more adsorbable state), whereas the iron phase increases the arsenic uptake capacity and reduces the band gap energy of photo-oxidation so that it can proceed under visible light instead of UV wavelengths.6 In combination, this results in rapid and extensive removal of contaminants like arsenic from solution in two steps, by the application of a single solid suspension.
qe =
Phosphate Analysis and UV−Vis Spectroscopy
Phosphate concentrations are very easily determined using UV−vis spectroscopy with the “molybdenum-blue” method, which is based on the formation molybdenum−phosphate with subsequent reduction using hydrazine sulfate.14 It is notable that recent work on the analytical chemistry of arsenate (a similar problematic oxy-anion) also focused on the development of UV−vis detection methods, and here the instructor could draw attention to the similar aqueous chemistries of phosphate and arsenate if this is considered relevant to the class.15 Calibration standards and analysis samples are run on the UV−vis spectrometer at an absorbance of 830 nm. A linear plot of absorbance versus known phosphate concentration can be used to determine the sample concentration by back-calculating from a calibration trend line. Concentrations of samples are measured, and the adsorption is compared to a calibration curve, that is, a welldefined linear plot of absorbance versus known solution concentrations. A calibration curve is constructed by plotting the absorbance versus the phosphate concentration. The best fit line through the data points is found with the method of leastsquares16 or, alternatively, with the Excel “add trend line” tool. Both methods allow for the calculation of the slope and intercepts of the best-fit line through the data points; however, they do not compute the uncertainties associated with these parameters. To estimate the uncertainties in the slope and intercept parameters an uncertainty analysis must be performed and it is important that the student learns this. We have provided error calculation notes as Supporting Information that the reader can incorporate into the course as they see fit.
The capacity of adsorbents for successfully removing the phosphate oxy-anion is determined using adsorption isotherms.12 The Langmuir (eq 1) and the Freundlich (eq 2) models describe the experimental data well. The models are given by
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GOALS AND ORGANIZATION OF THE LABORATORY MODULE The aim of the laboratory module is not only to learn the fundamental steps associated with the development of sorbents for the removal of toxic oxy-anions but also to provide the student with an opportunity to participate in the research process in the course of two to three sessions. The laboratory course requires approximately 10 h over two or three sessions, deliberately set apart to allow for sample preparation, with a typical attendance of 20 to 30 students. However the experimental procedure is quite flexible, allowing the classes to be condensed or extended in a variety of ways. The teaching assistant will need to be able to supervise ethanol evaporation and the use of a furnace outside of the taught session. Prior to the laboratory course, the student should have had classroom exposure to the introductory concepts of aqueous chemistry. After the experiments have been completed, all data handling can be completed using Microsoft Excel (or more advanced software if this is beneficial), allowing the students to
Q 0bCe 1 + bCe
(1)
qe = KFCe(1/ n)
(2)
(3)
where qe is the adsorbed concentration (in mg per g sorbent), V is the volume of solution used (in L), C0 is the initial concentration of sorbate in solution (in mg L−1), Ce is the sorbate concentration measured after sampling (mg L−1), and m is the dry mass of the sorbent used. If the sample is taken at equilibrium (i.e., when the adsorption reaction has continued until the sorbent reaches equilibrium), then Ce is the equilibrium concentration of sorbate in solution.
Adsorption Isotherms
qe =
V (C0 − Ce) m
respectively, where qe is the amount of adsorbate adsorbed at equilibrium per mass of sorbent, Ce is the concentration of adsorbate in solution, 1/n is a parameter reflecting the intensity of adsorption, Q0 is the maximum adsorption capacity at saturation, and KF and b are the Freundlich and Langmuir constants, respectively, determined experimentally by the student. The difference of the two models is that the Freundlich model assumes that the adsorbent can take up an infinite amount of adsorbate, albeit with weaker bonds toward the exterior of a multilayered system extending from the surface, whereas the Langmuir model assumes a monolayer adsorption that is saturated as soon as each site receives an ion.13 To calculate the amount of adsorbate adsorbed per gram of sorbent, a simple calculation can be used: 507
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9. The buffering systems used to maintain pH were sodium acetate/acetic acid for pH 5 and disodium tetraborate/dilute HCl/NaOH for pH 9. These buffering agents have a negligible interaction with the adsorption reactions under investigation. To produce phosphate concentrations of 25, 50, 75, and 100 mg L−1 respectively, 4 × 100 mL of pH 5 and 4 × 100 mL of pH 9 buffer solutions are measured into eight 100 mL plastic bottles. A 100 mL measuring cylinder should be used. Then a 0−5000 μL Eppendorf pipet is used to extract 2.5, 5.0, 7.5, and 10.0 mL, respectively, of buffer solution from each of the 100 mL solutions before filling the bottles. Next, 2.5, 5.0, 7.5, and 10.0 mL, respectively, of premade phosphate stock solution is transferred to the bottles (using the same pipet) to make up to 100 mL in all eight bottles. When pipetting, as few repetitions should be made as possible, by changing between 2.5 and 5.0 mL capacities. The bottles are then inverted for mixing, and magnetic stirrers are added to each. The pH is measured in every bottle for control. Next, one 3 mL standard is extracted from each bottle with a pipet and transferred to capped test tubes. A further 10 μL of every standard is transferred with a pipet to 8 additional capped test tubes, and 0.4 mL of hydrazine sulfate and 1 mL of Na2MoO4·2H2O are added to each. Distilled water is then used to top up the solutions to 11 mL each. This is the experimental control. The control is then placed in a water bath at 60 °C for 15 min. Meanwhile, 100 mg of the iron−titanium oxide bicomposite is measured into 8 numbered weighing boats. Each boat is emptied into the eight reaction bottles in order of 25 mg P L−1 up to 100 mg P L−1 (for pH 5 and then pH 9 in turn) by folding into a tube and flicking to ensure near-complete transfer of the bicomposite. The bottles are inverted twice and left on a magnetic stirrer to allow sufficient time (at least 24 h) for adsorption equilibrium. After adsorption equilibrium is achieved, 3 mL samples are extracted from each bottle and filtered using syringe filters. The resulting samples will then undergo UV−vis absorption spectroscopy to give the phosphate concentrations necessary for the calculation of adsorption isotherms.
learn good spreadsheet organization, graph plotting, regression fitting, and the use of formulas and functions to explore their results. We propose the following arrangements: Session 1: Perform the sorbent synthesis experiments and preparation of stock solutions (4 h). Session 2: Perform the adsorption experiments (2 h). Session 3: Determine concentrations from UV−vis absorbance data and data analysis (3 h). If student groups are larger, the UV−vis spectroscopic studies might be performed in parallel with the adsorption experiments (as samples are collected). This would condense the classes into two laboratory sessions if needed, with data analysis making an excellent coursework project to supplement.
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EXPERIMENTS
Mineral Synthesis (Session 1)
This synthesis procedure will produce ∼2.5 g of the bicomposite sorbent. Scaling up the quantities will result in a longer ethanol evaporation time, so it is recommended that instead each group of students prepares a batch of ∼8 sorbent preparations, each as detailed below (depending on the number of students involved). Ethanol, 200 mL, is transferred into a clean 200 mL volumetric flask, and a 0.6 M solution of Fe(NO3)3 is prepared by dissolving 48.48 g of Fe(NO3)3· 9H2O. The flask is warmed to room temperature under warm water. Then 30 mL of the reaction mixture is extracted and transferred into a 100 mL beaker. Then 1.50 g of titanium dioxide powder (anatase) is added to the beaker. The beaker is covered with wax film, and the contents are mixed on a magnetic stirrer for 30 min (or until homogeneous). The mixture is then sonicated for 30 min. The ethanol is evaporated, and the residue is calcined in a furnace for 10 min at 300 °C and then allowed to cool. The solid is crushed using a pestle and mortar to a fine powder and returned to the furnace at 300 °C for up to 6 h (or as long as time permits during the class). Finally the sorbent product is transferred to a clean, dry container for storage and session 2.
Phosphate Analysis Using UV−Vis Spectroscopy (Session 3)
Preparing the Stock Solutions (Phosphate and for UV−Vis Analysis)
A calibration curve is constructed using 10 mL solutions of standards at 0, 25, 50, 100, and 200 μg L−1 (which is equivalent to parts per billion, ppb). To each standard, 1 mL of sodium molybdate solution and 0.4 mL of hydrazinium sulfate are added, before they are made up to 10 mL volume with deionized water. Standards are measured using the UV−vis spectrophotometer at a wavelength of 830 nm; the absorbance is then recorded and plotted against phosphate concentration to establish a calibration curve (see Figure 2). The corrected absorbance is calculated by subtracting the average absorbance of the blank standard from each measured absorbance. The absorbance of the blank solution arises from the color of the starting reagents, reactions with impurities and reactions of interfering species. By correcting the absorbance values these effects are eliminated from the analysis.
The phosphate stock solution is prepared by dissolving potassium dihydrogen phosphate (KH2PO4) in deionized water to produce a stock solution of phosphate at a concentration of 1 g L−1 (i.e., 1000 ppm or μg/mL). For the purpose of analyzing the standards for phosphate concentration using the UV−vis spectrometer, a 500 mL stock solution of 25 g L−1 sodium molybdate is prepared by dissolving sodium molybdate (Na2MoO4·2H2O) in 5 M sulfuric acid. The correct amount of sodium molybdate is first dissolved in a small amount of 5 M sulfuric acid. Note that sulfuric acid is a hazardous acid to handle and the course administrator may choose to prepare this in advance. In addition, a 500 mL solution of 1.5 g L−1 hydrazine sulfate solution is prepared by dissolving hydrazine sulfate (NH2NH2·H2SO4) in deionized water.
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HAZARDS Students should wear laboratory coats, nitrile gloves, and eye protection and behave sensibly. Hazard assessments should be completed, understood, and signed by each participating student before the experiment is started. Students must understand that iron nitrate, titanium dioxide, and sodium molybdate can irritate the skin, eyes, and respiratory tract on
Adsorption Experiments (Session 2)
All equipment detailed below is washed with deionized water before usage, the pH meter is calibrated, and pipet tips are changed whenever used for a new solution. For the adsorption experiments, eight solutions are used: four different concentrations at pH 5 and four different concentrations again at pH 508
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Figure 2. Calibration curve for the UV−vis measurements, showing the absorbance increasing with phosphate concentration in μg L−1 (ppb). One standard solution (with 50 μg L−1) was omitted. Uncertainty reflects the differences between repeated measurements of the standard solutions.
Figure 3. The adsorption isotherms of P onto the bicomposite at pH 5 and 9 and the Langmuir and Frendlich model curves taken from one of the students’ data sets.
students need to plot the experimental data in linearized forms. For these isotherm models, linearized forms can be written as follows:
exposure. Hydrazine sulfate is a strong irritant, and sulfuric acid is corrosive. Students must be supervised at all times throughout the experiment and have access to suitable washes and eye baths. The class must be shown how to use the furnace correctly, with appropriate heat-protective gloves and tongs. Students should collect their waste on their workbenches, and the course administrator should arrange for proper disposal.
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Langmuir:
1 1 1 = + qe Q 0bCe Q0
Freundlich: ln qe = ln KF +
RESULTS AND DISCUSSION
⎛1⎞ ⎜ ⎟ln C e ⎝n⎠
With respect the Freundlich model, the slope and the intercept correspond to 1/n and ln KF. With respect to the Langmuir model, the slope corresponds to 1/bQ0 and the intercept to 1/ Q0. Using the data set displayed in Figure 3 and given in spreadsheet 2 as produced by one group of our students, we find for the Freundlich model at pH = 5 a regression line of y = 0.23x + 1.57, thus 1/n = 0.23, ln KF = 1.57 and KF = 4.8, and at pH = 9, a regression line of y = 0.25x + 1.02, thus 1/n = 0.25, ln KF = 1.02 and KF = 2.8. For the Langmuir model, we find at pH = 5, the intercept is 0.084, Q0 = 11.9, and the slope is 0.38 (=1/ bQb), therefore b = 0.22, and at pH = 9, the intercept is 0.14, Q0 = 7.4 and the slope is 0.66 (=1/Q0b), therefore b = 0.20. The values for the constants of the Langmuir model are summarized in Table 1. The student can derive the values for the slope and intercept again using the “add trend line” tool after plotting the graph or using the function option (i.e., SLOPE, INTERCEPT). Spreadsheet 2 shows all relevant calculations.
Phosphate Analysis and UV−Vis Spectroscopy
Here we are using the data set obtained by a group of two students working together (see spreadsheet 1 in the Supporting Information) during the class term in 2011−2012. The blank corrected absorbance data for calibration solutions are then plotted against the concentration of phosphate in the standards. If the “add trend line” tool of Excel is used to determine the linear regression line, then the calibration curve shown in Figure 2 is achieved with the corresponding equation, y = 0.0008x − 0.0024. Note that the slope and intercept values are derived in spreadsheet 1 graphically using the trend line option and using the function option in Excel. The equations outlined in the background and theory sections are then used to calculate the errors of the slope and the intercept (see spreadsheet 1 in the Supporting Information for calculations). It is good to encourage the students to set up their own spreadsheets. The uncertainties of the slope and intercept (one standard deviation, ±1 sd) are calculated as 0.0008 ± 0.0002 and −0.0024 ± 0.002, respectively.
Table 1. The Parameters for the Freundlich and the Langmuir Models at Different pH
Adsorption Isotherms
The adsorption isotherms for phosphate onto the bicomposite at the two different pH values (5 and 9) are shown in Figure 3 along with the two calibrated models, the Freundlich and the Langmuir equations (see spreadsheet 2 in the Supporting Information for calculations). Note that it would have been good to sample one aliquot between 0 and 5 μg mL−1; the students discussed this in their final report. To determine the model parameters (1/n and KF for the Freundlich model and Q0 and b for the Langmuir model), the
Model Parametersa Langmuir
Freundlich
pH
Q0
R2
b
1/n
KF
R2
5 9
11.9 7.4
0.93 0.99
0.22 0.20
0.23 0.25
4.8 2.8
0.90 0.97
a
See the spreadsheet 2 in the Supporting Information for the equations and codes.
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Now that the constants have been determined, the student is encouraged to compare the different models and the effect of changing pH and to discuss this with respect to surface chemical processes. We find that the uptake capacity is greater in acidic conditions and diminished in alkaline conditions; this is because the surface charge of the adsorbent becomes more positive as pH decreases, which is favorable for the adsorption of anionic species.7 Interested students are referred to D’Arcy et al. for further information.7 We note that the isotherm model parameters given in D’Arcy et al. are very slightly different from those presented in Table 1. This is because the sorbent was crushed more thoroughly and had a greater surface area; however, the relative differences in the parameters (e.g., higher uptake at lower pH) are the same.
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS The authors acknowledge financial support from Imperial College London, Lorraine Craig for organizational support, and the enthusiastic students of the different years. We thank the reviewers for their constructive comments.
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REFERENCES
(1) Anderson, D. A.; Gilbert, P. M.; Burkholder, J. M. Estuaries 2002, 25, 562−584. (2) Benner, S. Nat. Geosci. 2010, 3, 5−6. (3) Katayev, E. A.; Ustynyuk, Y. A.; Sessler, J. L. Coord. Chem. Rev. 2006, 250, 3004−3037. (4) Ali, I.; Gupta, V. K. Nat. Protoc. 2007, 1, 2661−2667. (5) Yang, R. T. Adsorbents - fundamentals and applications; Wiley Interscience: London, New York, 2003. (6) Zhou, W.; Fu, H.; Pan, K.; Tian, C.; Qu, Y.; Lu, P.; Sun, C.-C. J. Phys. Chem. C 2008, 112, 19584−19589. (7) D’Arcy, M.; Weiss, D. J.; Bluck, M.; Vilar, R. J. Colloid Interface Sci. 2011, 364, 205−212. (8) Antelo, J.; Avena, M.; Fiol, S.; López, R.; Arce, F. J. Colloid Interface Sci. 2005, 285, 476−486. (9) Bang, S.; Patel, M.; Lippincott, L.; Meng, X. Chemosphere 2005, 60, 389−397. (10) Schuttlefield, J. D.; Larsen, S. C.; Grassian, V. H. J. Chem. Educ. 2008, 85, 282−284. (11) Tribe, L.; Barja, B. C. J. Chem. Educ. 2004, 81, 1624−1627. (12) Limousin, G.; Gaudet, J. P.; Charlet, L. Appl. Geochem. 2007, 22, 249−275. (13) McCash, E. M. Surface Chemistry; Oxford University Press: Oxford, 2001. (14) Kharat, S. J.; Pagar, S. D. E-J. Chem. 2009, 6, S515−S521. (15) Morita, K.; Kaneko, E. Anal. Sci. 2006, 22, 1085−1090. (16) Harris, D. C. Quantitative Chemical Analysis; 8th ed.; W. H. Freemann & Co.: 2010. (17) The Freundlich and Langmuir isotherms are usually considered as purely empirical equations, which does not allow for a description and interpretation of the coefficients. In order to overcome this difficulty, Skopp derived the Freundlich and the Langmuir adsorption isotherms from fractal kinetics. 18 This provides a physical interpretation of the coefficients in the equations in terms of the molecular characteristics of the adsorbent surface and the adsorbate.18 (18) Skopp, J. J. Chem. Educ. 2009, 86, 1341−1343.
CONCLUSIONS Students are able to safely and inexpensively synthesize a mixed oxide bicomposite sorbent using a variety of laboratory techniques. The product is then tested for its oxy-anion uptake capacity in acidic and alkaline conditions, using the isotherm approach. The data collected allows two different isotherm models to be compared, the Langmuir and Freundlich models for solute adsorption onto a solid. The Langmuir model better fits the data compared to the Freundlich model, and thus it should be preferred when modeling the adsorption of phosphate and similar chemical species onto a solid, such as iron oxide and titanium oxide. The experiment illustrates how a simple series of calculations using two different models could lead to an interesting and insightful representation of the adsorption process. This experiment is particularly relevant because its results can be used to make predictions on the dispersal and partitioning of elements in soils (iron oxide being a major constituent). In particular, this technique is extremely useful to map out and determine the amounts of hazardous pollutants in soils. These calculations, although relatively simple, are essential for determining the amount of target oxy-anion adsorbed, and they can be used as a guide to determine the cleanup option required. Further applications include the use of this adsorbent in wastewater treatment processes. These calculations allow one to estimate how much phosphorus would be removed from water treated with this adsorbent. It is essential to control the concentration of phosphate in discharge water because excessive concentrations may lead to eutrophication in the natural environment. For example, consider a phosphorus mine where acidic mine water, with a phosphorus concentration of 4 mg L−1, is produced. According to the Langmuir model built with the experimental data for pH 5, about 5.7 mg of phosphate would be adsorbed per gram of adsorbent introduced in the mine water. This value would increase to 6.8 mg of phosphate per gram of adsorbent if the Freundlich model were applicable.17 If the mine water were to be discharged into a stream, treatment with the Fe2O3−TiO2 adsorbent would be sufficient, as most regulations require a concentration of P in solution lower than 1 mg L−1: achievable with this adsorbent.
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ASSOCIATED CONTENT
S Supporting Information *
Student notes, instructor notes, and spreadsheets. This material is available via the Internet at http://pubs.acs.org. 510
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