Environ. Sci. Technol. 1989, 23,580-587 Dadswell, M. J. N a t . Mus. N a t . Hist. Ottawa, Published in Zoology 11, 1974. Edmondson, W. T., Ed. Fresh Water Biology, 2nd ed.; John Wiley and Sons: New York, 1957. Pennak, R. W. Freshwater Invertebrates of the United States; Ronald Press: New York, 1953. Scott, W. B.; Crossman, E. J. Bull.-Fish. Res. Board Can. 1973, No. 184. Nilssen, J. P. Rep.-Znst. Freshwater Res. Drottningholm 1984, NO. 61, 138-147. U.S. EPA The Acidic Deposition Phenomenon and its Effects. EPA-600/8-83-016B,1985. Eadie, M. McA.; Keast, A. Can. J. 2001.1984,62,1689-1695. Schindler, D. W. In The Earth i n Transition: Patterns and Processes of Biotic Impoverishment; Woodwell, G. M., Ed.; Cambridge University Press: New York, in press. Schindler, D. W.; Kasian, S.E. M.; Hesslein, R. H. Environ. Monit. Assess., in press. Gorham, E.; Martin, F. B.; Litzau, J. T. Science (Washington, D.C.) 1985,225, 407-409. Kramer, J. R.; Andren, A. W.; Smith, R. A.; Alexander, R. B.; Oehlert, G. In Acid Deposition: Long-Term Trends; National Academy Press: Washington, DC, 1986; Chapter 7, p p 231-299. Schindler, D. W. Limnol. Oceanogr., in press. Henrickson, L.; Oscarson, H. G.; Stenson, J. A. E. In Ecological Impact of Acid Precipitation; SNSF Project: Oslo, Norway, 1980; pp 316-317. Rapport, D. L.; Regier, H. A,; Hutchinson, T. C. Am. Nat. 1985,125, 617-640.
Rogalla, J. A.; Brezonik, P. L.; Glass, G. E. Water, Air, Soil Pollut. 1986, 31, 95-100. Kratz, T. K.; Cook, R. B.; Bowser, C. J.; Brezonik, P. L. Can. J . Fish. Aquat. Sci. 1987, 44, 1082-1088. Davis, R. B.; Anderson, D. S.; Berge, F. Nature (London) 1985, 316,436-438. Effler, S. W.; Schafran, G. C.; Driscoll, C. T. Can. J . Fish. Aquat. Sci. 1985,42, 1707-1711. Wiener, J. G.; Eilers, J. M. J . Lake Reserv. Manage. 1987, 3, 365-378. Rago, P. J.; Wiener, J. G. Trans. Am. Fish. SOC.1986,115, 438-447. Rahel, F. J. Can. J. Fish. Aquat. Sci. 1986, 43, 124-134. Rahel, F. J.; Magnuson, J. J. Can. J. Fish. Aquat. Sci. 1983, 40, 3-9. Minns, C. K. In Environmental Impact Models to Assess Regional Acidification; Reidel Publishing Co.: Dordrecht, The Netherlands, in press. Elton, C. S. T h e Ecology of Invasions by Animals and Plants; Metheun: London, 1958. Pimm, S.L. Nature (London) 1984, 307, 321-326. Schindler, D. W. Can. J . Fish. Aquat. Sci. 1987,44, Suppl. 1, 26-33. Rudd, J. W. M.; Kelly, C. A.; Schindler, D. W.; Turner, M. A. Science (Washington, D.C.) 1988,240, 1515-1517. Schindler, D. W. Oikos, in press. Received for review November 17,1987. Accepted November 28, 1988. Financial support for the work was provided by the Canadian Department of Fisheries and Oceans.
Calorimetric Investigation of Anion Adsorption onto Goethite Mlchael L. Machesky," Brian L. Blschoff, and Marc A. Anderson
Water Chemistry Program, University of Wisconsin, Madison, Wisconsin 53706 Titration calorimetry was employed to investigate adsorption of fluoride, iodate, phosphate, and salicylate anions onto goethite at constant temperature (25 "C), pH (4.0), ionic strength (0.05 M), and solid concentration (10 g/L). In addition, adsorption data were collected at three temperatures (10,25, and 40 "C) and proton stoichiometry values (H+ adsorbed/anion adsorbed) were monitored as well. Measured adsorption enthalpies were exothermic and decreased with increasing surface coverage for all anions. Furthermore, the enthalpy curves all exhibited the same general shape. That is, at low surface coverages (20% and 70%) enthalpies again decreased more rapidly and even became slightly endothermic for phosphate and fluoride. The primary reason for this similarity in shape is believed to be the surface site distribution inherent to the goethite. Except for iodate, calculated enthalpies did not agree with those measured. It is also postulated that anion adsorption free energies have a large favorable entropic component with enthalpic factors of more minor importance. Introduction Anion adsorption influences several processes of environmental concern including contaminant transport, nu-
* Present address: Department of Geosciences, 208 Deike Building, The Pennsylvania State University, University Park, PA 16802. 580
Environ. Sci. Technol., Vol. 23,No. 5, 1989
trient availability, and mineral dissolution rates (1-3). Much of our current state of understanding concerning adsorption processes has resulted from laboratory studies dealing with well-characterized single-component, adsorbent-adsorbate systems, typically a hydrous metal oxide and a single cation or anion. Under these conditions, results are usually successfully interpretable by any of several available adsorption models ( 4 ) . However, these models have been less successful when predicting competitive adsorption effects (5),especially when predictions are based on single adsorbent-adsorbate experimental results (6). This suggests quantitative assessment of the importance of adsorption phenomena in the field is not yet possible and, on a more fundamental level, that the mechanisms underlying adsorption processes are incompletely understood. This lack of mechanistic understanding exists primarily because solid-solution partitioning experiments (often as a function of pH) have been the predominant method used to infer adsorption mechanisms. In order to increase our mechanistic understanding of adsorption processes, new or underutilized experimental techniques must be applied to obtain a wider variety of data. Recently, this has begun to occur as investigations employing pressure jump and electric field pulse kinetic techniques (7,8),direct potential measurements with metal oxide electrodes (9, IO), and in situ spectroscopic methodologies such as electron spin resonance (11,12),cylindrical internal reflectance-FTIR (13, 14), and synchrotron radiation studies (15) have appeared. From such studies it appears that anion adsorption onto hydrous metal oxides
0013-936X/89/0923-0580$01.50/0
0 1989 American Chemical Society
Table I. Anion Titrant Compositions, Amounts Added, Analysis Methods, and Detection Limits
anion
molarity (compositn)
0.40 M (16.8 g of NaF/L + 100 mL of 0.4M HF) 1030.10 M (19.79 g of NaIOB/L+ 100 pL of 1 M HNOB) 0.20 M (27.6 g of NaH2P04.H20/L+ 120 pL of 5 M NaOH) H2P04C6H40HC00- 0.10 M (16.01 g of C8H,0HCOONa/L + lo00 pL of 1 M HNOq)
F
can result in inner-sphere or outer-sphere surface complexation, and a two-step ligand-exchange mechanism is suggested based on rapid kinetic studies (16, 17):
+ H+ = SOH2+ SOH2++ A- = SA + H 2 0 SOH
site activation
(1)
ligand exchange (2)
where SOH represents a surface hydroxyl group and Aan anion. The site activation step accounts for the rapid pH increase that is frequently observed during anion adsorption studies (18). According to this mechanism, the amount of acid required to return to the pH value before anion addition equals the number of surface groups protonated during the site activation step. Also, this amount of acid, divided by the quantity of anion adsorbed, is termed the adsorption, exchange, or proton stoichiometry ratio (19, 20). This ratio reflects charge compensation processes in the inner (Stern) part of the double layer (21Z and it has also been used in adsorption equilibrium constant expressions and to help interpret adsorption mechanisms for cations (pH decreases so back-titration with base is used) as well as anions (19, 22). Titration calorimetry, with which the heat changes accompanying adsorption processes can be determined, has been a relatively underutilized method. The technique has been extensively used to study homogeneous solution reactions (23),but relatively few studies involving solidaqueous solution systems have appeared. These include studies of cation-exchange processes (24,25),phosphate adsorption onto allophanic soils (26)and, more recently, several studies involving hydrous metal oxides (14,27,28). Titration calorimetry is a useful technique to study adsorption reactions for several reasons. First, adsorption reaction enthalpies are obtained, and the enthalpy of a reaction (not the free energy) is a 'direct measure of the bond strength between associating species. Second, if free energy data are also available, the entropy of adsorption reactions can be calculated and consequently a complete thermodynamic description of adsorption processes can be obtained. Third, adsorption enthalpies can be used to predict how temperature changes should influence adsorption processes at equilibrium and temperature has not been considered an important variable in most previous adsorption studies. Those studies that have considered temperature a variable, however, have most often found that as temperature increases, anion adsorption decreases (29,30), while cation adsorption increases (31). Finally, titration calorimetry is a potentially powerful method for determining the distribution of surface site energies that exist on hydrous metal oxide surfaces since analogous calorimetric studies of gas adsorption onto solids have been used extensively for this purpose (32). In this investigation, titration calorimetry was used to study anion adsorption onto goethite. Goethite is an ideal adsorbent for several reasons. First, it is the most common iron oxide mineral at the earths surface ( 1 7). Second, its surface chemistry is fairly well understood, and in particular, the distribution and abundance of surface hydroxyl groups participating in adsorption reactions are known (2,
anion added/ increment, pmol
anal. method
160 40 50 15
specific ion electrode (38) colorimetric (39) colorimetric (40) absorbance at 296 nm
pract detectn limit, p M 1 2 0.3 2
17). Finally, in our laboratories, goethite has been the primary substrate for several related studies which have dealt with charge development (33),calorimetric (34)and in situ infrared spectroscopic (35)characterization of its surface properties, and adsorption behavior. For this study, ionic strength, pH, and goethite suspension concentrations were held constant, thereby reducing the experimental variables to (1) the nature of the anion adsorbed and (2) the extent of surface coverage.
Experimental Section Goethite was prepared according to a method of Atkinson et al. (36)and is described in more detail elsewhere (27, 34). X-ray diffraction confirmed the solid was goethite, although diffraction peaks were broadened due to the small crystallite size (about 50 X 15 X 10 nm from TEM micrographs). The N2 BET surface area was 81 m2/g. Suspensions were prepared as needed from freeze-dried solid. Goethite (4.5 g) was added to 450 mL of a 0.05 M NaN03 solution [prepared with ultrapure (>99.996%) NaN03 and Milli-Q H20]contained in a 500-mL polycarbonate centrifuge bottle. The suspension was acidified with 3 mL of 0.2 N HN03 (ultrex grade) to pH -4.5. This was followed by -24 h of intermittent ultrasonic dispersion over a 1-week period, during which time more 0.2 N HN03 was added (- 1mL total) to keep the pH between 4.2 and 4.5. Any particular suspension preparation was exhausted in less than 1 month to minimize any long-term chemical, physical, or biological aging processes. The apparent zero point of charge (zpc) for these air-equilibrated suspensions was determined (by potentiometric titrations) to be 8.0 f 0.2, for which an independent confirmation (8.1 f 0.2) exists (33, 37). The anions investigated during this study, salicylate (C6H40HC00-),phosphate (H2P04-),iodate (IO3-), and fluoride (F-), were chosen because they differed in size, their solution chemistry was relatively uncomplicated at pH 4 (i.e., distance from dissociation constants), and simple and sensitive analysis methods existed for their detection. Anion titrant solutions (Table I) were made with ACS reagent-grade (or better) chemicals and were stored at 6 "C. The acid or base added to these solutions brought pH values to between 4 and 5, which minimized changes in solution blank pH values. Titrant concentrations were verified with the primary standards recommended for the analysis procedures (literature references given in Table I). Adsorption experiments employed a wide-mouth 250mL polymethylpentene (PMP) container jacketed with tightly coiled copper tubing. A water bath circulated water through the tubing to maintain the desired temperature (10,25, or 40 "C). Holes drilled through the lid of the PMP container accommodated miniature pH electrodes (Microelectrodes, Inc., Londonderry, NH), which were standardized with 1:l phosphate and 0.05 M phthalate buffers at the desired temperature, a capillary tube for dispensing acid, and an Eppendorf or volumetric pipet tip. Goethite suspension (200 mL) was added to the PMP Environ. Sci. Technol., Vol. 23,No. 5, 1989
581
container and 0.5-1 h allowed for thermal equilibration at pH 4.00 f 0.02 (with magnetic stirring). Next, an aliquot of anion titrant was added with an Eppendorf pipet and the pH allowed to drift freely for 1 min. If the pH increased more than 0.03 pH unit, standardized 0.2 N HNO, was dispensed from a Gilmont precision microburet (through the capillary tubing) to return the pH to 4.00 and the volume of added acid recorded. After an additional 4 min of equilibration, a 10-mL aliquot of the suspension was withdrawn with a volumetric pipet and immediately filtered through a 0.05-pm Nuclepore polycarbonate filter using a pneumatic pressure filtration apparatus (Schleicher & Schuell, Inc., Keene, NH) also jacketed with copper tubing to maintain the desired temperature during filtration. The first milliliter of filtrate was discarded, and filtration took 15-40 min depending on the degree of aggregation induced by the added anion. Eleven equalvolume additions were made for each anion at each temperature, and solution blanks were treated similarly except filtration was unnecessary. Goethite dissolution was minimal, since in representative filtrates from the anion adsorption experiments dissolved Fe was always -5
a
-I
20% and 70%) adsorption enthalpies decrease more rapidly and even appear to become slightly endothermic for phosphate and fluoride. The precision of the data for salicylate at surface coverages of >20% is very poor because the measured heat changes were close to the practical detection limit of the calorimeter (34).
Discussion Adsorption data were collected at three temperatures so adsorption enthalpies could be calculated and compared to those measured. A form of the Clausius-Clapeyron equation can be used for this purpose (30):
$-)
4?/( $-
AHI = R In C1
(5)
where AHI is the isosteric heat of adsorption at a given surface coverage or adsorption density, R is the gas constant, C2is the equilibrium solution concentration of the ion at temperature T2and the given surface coverage, and C1 is the equilibrium concentration of the ion at temperature T1and the given surface coverage. The amount of an ion adsorbed at a given adsorption density must decrease with increasing temperature for the calculated enthalpy to be exothermic. This presupposes that the enthalpy is constant over the temperature range considered and several factors can invalidate this assumption, including a nonzero heat capacity and shifts in solution- or Environ. Sci. Technol., Vol. 23, No. 5, 1989 583
Table 11. Residual Solution Concentrations (pM)and Amounts Adsorbed (pmol/g) from the 25, 10, and 40 OC Adsorption Experiments
-iodate
fluoride
phosphate
soln
ads
soh
ads
7 14 34 44 62 116 162 275 401 654 991
20 40 60 83 106 128 152 171 191 201 205
30 97 247 521 985 1585 2359 3293 4387 5571 6389
78 156 230 297 349 394 427 452 469 490 509
6 14 24 34 56 93 149 228 366 578 887
20 40 61 84 107 130 153 176 194 209 216
34 101 243 524 996 1601 2389 3337 4365 5582 6935
77 155 230 296 348 392 425 448 413 488 501
23 32 47 59 100 137 194 317 430 547 819
18 38 59 81 102 126 148 166 188 212 223
46 112 258 504 941 1528 2261 3170 4229 5343 6371
76 154 229 298 354 400 437 465 485 511 519
salicylate
soln
ads
soln
ads
0 0 0 0 6 60 245 535 896 1262 1735
25 52 80 109 140 168 185 194 199 207 208
2 6 26 56 97 145 204 264 339 434 514
8 15 21 27 32 38 43 48 53 51 64
0 0 0 2 15 111 307 566 930 1283 1730
25 52 80 109 139 163 179 191 196 205 208
2 4 27 60 100 146 201 259 334 413 508
8 15 21 27 32 38 43 48 53 59 64
0 0 0 0 2 38 222 510 901 1325 1847
25 52 80 109 141 171 191 203 208 214 215
2 8 27 53 90 141 194 260 339 433 521
7 15 21 27 33 38 43 48 53 57 63
25 "C
10 "C
40
solid-phase equilibria (e.g., species distributions, zero points of charge) that could cause the adsorption mechanism or the magnitude of the adsorption enthalpy to vary with temperature. Except for iodate, the measured exothermic adsorption enthalpies were not consistent with those calculated by eq 5. There are several possible reasons for this discrepancy. First, analytical constraints coupled with the observed decrease in adsorption enthalpy with increasing surface coverage may have obscured differences in residual solution concentrations. For fluoride, calculated adsorption enthalpies are exothermic for the first three entries from Table I1 at 40 and 25 "C (AHa&= -23, -8, and -2 kJ/mol, respectively), and the measured adsorption enthalpies are also only slightly exothermic. This implies that only small differences in residual solution concentrations with temperature are to be expected. For example, the difference in residual solution concentrations between 40 and 25 "C is only +4% for the -2 kJ/mol calculation, which is less than the estimated error in the measured concentrations (*lo%). The salicylate (at surface coverages of >20%) and phosphate results were probably affected similarly. For phosphate, measured adsorption enthalpies are appreciably exothermic to 70% surface coverage but residual solution concentrations are 70%) surface coverages. In this region measured adsorption enthalpies appear to become slightly endothermic and therefore residual solution concentrations would be expected to decrease at higher temperatures. However, other factors may play a role at high surface coverages such as changes in the adsorption mechanism. A final factor may be that adsorption equilibrium was not established within the time frame of the experiments. Adsorption rates probably increase with temperature, and so for the short reaction times used in this study (
