Influence of Temperature on the Adsorption of Mellitic Acid onto Kaolinite

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Langmuir 2006, 22, 4208-4214

Influence of Temperature on the Adsorption of Mellitic Acid onto Kaolinite Michael J. Angove,* John D. Wells, and Bruce B. Johnson Colloid and EnVironmental Chemistry Laboratory, La Trobe UniVersity, P.O. Box 199, Bendigo, Victoria 3552, Australia ReceiVed December 21, 2005. In Final Form: February 23, 2006 The adsorption of mellitic acid (benzene-1,2,3,4,5,6-hexacarboxylic acid) onto kaolinite was investigated at five temperatures between 10 and 70 °C. Mellitic acid adsorption increased with increasing temperature at low pH (below pH 5.5), but at higher pH, the effect of increasing temperature was to reduce the amount adsorbed. Potentiometric titrations were conducted, adsorption isotherms were measured over the same temperature range, and the data obtained were used in conjunction with adsorption edge and ATR-FTIR spectroscopic data to develop an extended constant capacitance surface complexation model of mellitic acid adsorption. A single set of reactions was used to model all data at the five temperatures studied. The model indicates that mellitic acid sorbs via outer-sphere complexation to surface hydroxyl (SOH) groups on the kaolinite surface rather than to permanent charge sites. The reactions proposed are SOH + L6- + 2H+ h [(SOH2)+(LH)5-]4- and SOH + L6- h [(SOH)(L)6-].6- Thermodynamic parameters calculated from the temperature dependence of the equilibrium constants for these reactions indicate that the adsorption of mellitic acid onto kaolinite is accompanied by a large entropy increase.

Introduction Humic substances are known to sorb to mineral surfaces in soil and aquatic systems, causing a substantial change in the surface characteristics of the minerals.1,2 Hence, before we can hope to understand the nature of interactions between particulate matter and solution species in soil and sediment systems we must first understand the nature of the interaction between the minerals themselves and organic matter. There have been many investigations of the sorption of humic and fulvic acids and soil minerals in recent years which have provided useful insights into these complex systems.3-5 However, because fulvic and humic molecules are both large and complex, more detailed information on the chemical nature of the interaction can be obtained by studying the sorption behavior of relatively simple analogues of these macromolecules. The results of these experimental studies are often combined with spectroscopic data in order to deduce the most likely set of surface reactions generally by use of a surface complexation model.6-8 While the set of reactions proposed depends, at least to some extent, on the particular surface complexation model chosen, infrared spectroscopic information can provide direct evidence for the nature of the complex (whether it involves inner- or outersphere complexation) and on the extent of protonation of the surface complex when the adsorbing ligand contains carboxylate groups.9 * To whom correspondence should be addressed. E-mail: m.angove@ latrobe.edu.au. (1) Surface and Colloid Chemistry in Natural Waters and Water Treatment; Beckett, R., Ed.; Plenum: New York, 1990. (2) Marshall, S. J.; House, W. A.; Russell, N. J.; White, G. F. Colloids Surf., A 1998, 144, 127-137. (3) Wershaw, R. L.; Leenheer, J. A.; Sperline, R. P.; Song, Y.; Noll, L. A.; Melvin, R. L.; Rigatti, G. P. Colloids Surf., A 1995, 96, 93-104. (4) Vermeer, A. W. P.; van Riemsdijk, W. H.; Koopal, L. K. Langmuir 1998, 14, 2810-2819. (5) Yoon, T. H.; Johnson, S. B.; Brown, G. E. J. Langmuir 2004, 20, 56555658. (6) Boily, J.-F.; Persson, P.; Sjo¨berg, S. Geochim. Cosmochim. Acta 2000, 64, 3453-3470. (7) Lackovic, K.; Johnson, B. B.; Angove, M. J.; Wells, J. D. J. Colloid Interface Sci. 2003, 267, 49-59. (8) Johnson, S. B.; Yoon, T. H.; Kocar, B. D.; Brown, G. E. J. Langmuir 2004, 20, 4996-5006.

Although phyllosilicate minerals are very important in natural systems, there have been fewer studies of the adsorption of organic molecules to clays than to (hydr)oxide substrates, and those that are published have generally not involved concurrent spectroscopic investigation. This is not surprising as phyllosilicate minerals generally adsorb polar organic species less strongly than (hydr)oxide minerals, and consequently, the adsorption bands due to the adsorbed species are often weak. Where ATR-FTIR studies of carboxylic acid adsorption to clay minerals have been reported, these have included neither macroscopic adsorption measurements nor surface complexation modeling.10,11 The available spectroscopic evidence for the type of binding between organic acids and kaolinite provides no clear generalizations. Kubicki et al.10 studied the adsorption of acetic, oxalic, benzoic, citric, and salicylic acids on kaolinite and found no evidence of inner-sphere complexation. For most of these adsorbates, they found that simple washing with distilled water resulted in the ATR-FTIR spectra of the adsorbed species disappearing. In contrast, Specht and Frimmel11 proposed innersphere complexation of oxalic, malonic, and succinic acids to kaolinite based on changes in spectra on adsorption of these molecules. Temperature is an important variable in natural water and soil systems, with seasonal variations commonly exceeding 30 K. In regions with hot climates, surface soil temperatures often reach 50 °C in summer but fall below 5 °C in winter. Temperature variations are known to effect the uptake of nutrients by plants: for example, Schwartz et al.12 showed that variations of temperature in the root zone significantly affect the uptake of phosphate and zinc by barley. Since soil organic matter influences the availability of trace nutrients, changes in its sorption can influence the nutrient status of plants. (9) Boily, J.-F.; Nilsson, N.; Persson, P.; Sjo¨berg, S. Langmuir 2000, 16, 57195729. (10) Kubicki, J. D.; Schroeter, L. M.; Itoh, M. J.; Nguyen, B. N.; Apitz, S. E. Geochim. Cosmochim. Acta 1999, 63, 2709-2725. (11) Specht, C. H.; Frimmel, F. H. Phys. Chem. Chem. Phys. 2001, 3, 54445449. (12) Schwartz, S. M.; Welch, R. M.; Grunes, D. L.; Cary, E. E.; Norvell, W. A.; Gilbert, M. D.; Meredith, M. P.; Sanchirico, C. A. Soil Sci. Soc. Am. J. 1987, 51, 371-375.

10.1021/la0534571 CCC: $33.50 © 2006 American Chemical Society Published on Web 03/25/2006

Adsorption of Mellitic Acid onto Kaolinite

While the effect of temperature on the adsorption of metal ions to mineral surfaces has been the subject of several studies,13-17 there have been few studies of the effect of temperature on the adsorption of polar organic species to clay minerals. Those that are available are often studies on soil and sediment systems where the surface chemistry of the substrates is uncertain at best. An exception is the work of Ward and Brady,18 who investigated the effect of temperature on the adsorption of oxalate onto kaolinite and found that adsorption increased at the higher temperature below pH 6 but decreased above pH 6. The present study investigates the effect of temperature on the adsorption of mellitic acid (benzene hexacarboxylic acid) onto kaolinite at a range of temperatures between 10 and 70 °C, with the aim of elucidating the species present on the surface. It combines potentiometric titrations and adsorption measurements as a function of both pH and concentration with ATR-FTIR spectroscopy. We have used a simple extended constant capacitance surface complexation model (ECCM) to describe the sorption data at the various temperatures using surface species that are consistent with the spectroscopic measurements. This model, although simplistic in its description of electrostatics and surface properties, is capable of predicting mellitate sorption over a wide range of temperature, pH, and concentration conditions. The results from this study are compared with those of a previous study of the adsorption of mellitic acid to goethite as a function of temperature.19 Experimental Section Characterization of Kaolinite. Acid-washed kaolinite from Ajax Chemicals (Sydney, Australia) was used without further treatment. The BET surface area (measured with a Micromeritics ASAP 2000 instrument after degassing at 25 °C for 18 h) was 14.7 m2 g-1. XRD analysis showed that the kaolinite sample was free of mineral contaminants. Equipment and General Procedure. Adsorption experiments and potentiometric titrations were performed in a 300-mL jacketed reaction vessel under an atmosphere of CO2-free nitrogen. The reaction vessel and flat flange lid of borosilicate glass were separated by a neoprene rubber gasket. Temperature was maintained ((0.5 °C) by passing water through the jacket (20 L min-1) from a thermostated water bath. Experiments were performed at 10, 25, 40, 55, and 70 °C. A 2-m air condenser (5 mm diameter) minimized the loss of water by evaporation at the higher temperatures; at 70 °C, less than 2% of the water was lost from the vessel after 72 h, the longest period used for any titration experiment. Ross Sure-Flow double-junction combined electrodes (Orion Research) were used to measure pH. They were contained completely within the thermostated vessel and calibrated with NBS standard buffers at the reaction temperature. Electrode calibration was checked at the end of each experiment, and the results of the experiment discarded if the pH varied by more than 0.05 from that expected for the buffer at that temperature. The mass of kaolinite used was that required to give a surface concentration of 94 m2 L-1, and the supporting electrolyte was 5 mM KNO3.. The suspension was stirred using a PTFE-coated magnetic stirrer and was pre-equilibrated for 18-24 h under a nitrogen atmosphere at the natural suspension pH before the commencement of any experiment. A longer equilibration time (48 h) was allowed for experiments at 10 °C. All experiments were performed in triplicate, (13) Tewari, P. H.; Lee, W. J. J. Colloid Interface Sci. 1975, 52, 77-88. (14) Fokkink, L. G. H.; de Keizer, A.; Lyklema, J. J. Colloid Interface Sci. 1990, 135, 118-131. (15) Johnson, B. B. EnViron. Sci. Technol. 1990, 24, 112-118. (16) Goldberg, S.; Forster, H. S. Soil Sci. 1998, 163, 109-114. (17) Angove, M. J.; Johnson, B. B.; Wells, J. D. J. Colloid Interface Sci. 1998, 204, 93-103. (18) Ward, D. B.; Brady, P. V. Clays Clay Miner. 1998, 46, 453-465. (19) Angove, M. J.; Wells, J. D.; Johnson, B. B. J. Colloid Interface Sci. 2006, 296, 30-40.

Langmuir, Vol. 22, No. 9, 2006 4209 and all experimental data points are shown in the figures (with the exception of titrations, where, for clarity, only every third point from each titration is shown). Thus, each figure provides an indication of the magnitude of the variation (ca. 3-5%) in the experimental methods and data. Equilibration Times. Kinetic experiments were conducted to determine the period required for adsorption to reach equilibrium. An aliquot of adsorbate was added to a mineral suspension, and the pH varied in steps from 3 to 9. At each pH, samples were taken immediately and then again after 5, 10, 15, 20, 30, 60, and 120 min. Ligand concentrations in solution remained constant after approximately 10 min at all temperatures except 10 °C, where 15 min was required. Adsorption Edges. After pre-equilibration, the suspension was titrated to about pH 3 with 0.100 M HNO3, and an aliquot of stock 10 mM mellitic acid solution added to provide a total mellitic acid concentration of 100 µM. The pH was then increased in increments of 0.2-0.5 to about 10 by addition of 0.100-M KOH. An equilibration period of about 25 min was allowed between each increment, after which a 5 mL sample was taken from the suspension. Samples were filtered immediately through 0.22-µm polycarbonate membrane filters (Poretics), and the supernatant solutions analyzed by highperformance liquid chromatography (HPLC) (Shimadzu LC-10AI). A Vydac #302IC4.6 anion-exchange column was used with a 5 mM H2SO4 mobile phase. HPLC allowed the separation of nitrate, which has a broad UV signature, from mellitic acid. Detection was by photo-diode-array detector (Shimadzu SPD-M10AVP), with the greatest sensitivity achieved at 230 nm. The detection limit for mellitic acid by this method was ca. 1 µM. Adsorption of ligand to the membrane filter was small ( 0.17) at either pH. Modeling the Data and Determination of Thermodynamic Parameters. Many different adsorption reactions were tested. Those proposed represent the only set that was consistent with the spectroscopic data while fitting the titration and adsorption edge data at all temperatures with the minimum number of adjustable parameters. Mellitate adsorption onto kaolinite was best represented by a model that included two outer-sphere mellitate complexes. (Outer-sphere complexes are indicated by square brackets [ ].) The full set of model parameters at each temperature are shown in Table 1. At lower pH, the predominant reaction proposed in the model is

SOH + L6- + 2H+ h [(SOH2)+(LH)5-]4The value of the equilibrium constant for this reaction (Table 1) increased by about 1.5 orders of magnitude as temperature increased. The second reaction

SOH + L6- h [(SOH)(L)6-]6became more important at higher pH values. For this second reaction, the equilibrium constant did not change significantly with temperature. Table 3 shows the values obtained for the thermodynamic parameters from the temperature dependence of the equilibrium constants derived from the surface complexation modeling. These results indicate that the formation of [(SOH2)+(LH)5-]4- on kaolinite is an endothermic process, while the enthalpy for the formation of [(SOH)(L)6-]6- is indistinguishable from zero.

Table 3. Thermodynamic Parameters for Mellitic Acid Adsorption onto Kaolinite from Surface Complexation Modeling reaction

∆H (kJ mol-1)

∆S (J K-1 mol-1)

SOH + L6- + 2H+ h [(SOH2)+(LH)5-]4SOH + L6- h [(SOH)(L)6-]6-

50 ( 10 0(8

500 ( 30 80 ( 20

Adsorption entropies for both reactions are positive but are much larger for the formation of [(SOH2)+(LH)5-]4- than for [(SOH)(L)6-].6-

Discussion Effect of Temperature on Mellitic Acid Adsorption. Figure 3 indicates that adsorption increases with temperature at low pH but decreases with increasing temperature at high pH. This distinctive behavior has been noted previously for the oxalatekaolinite18 and mellitic acid-goethite19 systems. For the mellitic acid-kaolinite system, both the increase in adsorption at low pH and the decrease in adsorption at high pH can be explained by considering the electrostatic interactions between mellitate species and the surface sites. The equilibrium constants for the surface protonation reactions (Table 1) show that at higher pH, where protonation of the SOH sites is unlikely, an increase in temperature gives rise to increased deprotonation of both the SOH and the XH surface sites, resulting in a larger net negative surface charge at higher temperatures. The decrease in adsorption with increasing temperature at higher pH therefore most probably results from the increase in electrostatic repulsion between the negatively charged mellitate ion and the more negative surface charge. At low pH, most of the SOH sites are protonated. Table 1 indicates that the equilibrium concentration of SOH2+ sites increases with temperature, making the electrostatic interaction between these sites and the mellitate species in solution more favorable. Thus, the amount adsorbed increases as the temperature is increased. This increase in mellitic acid sorption with increasing temperature at low pH is much more pronounced on kaolinite than was found previously for goethite.19 The surface protonation constants for goethite indicate that the surface concentration of SOH2+ sites tends to decrease with increasing temperature at low pH, reducing the electrostatic interaction with mellitate species; Angove et al.19 attributed the slightly enhanced adsorption to goethite at low pH to reduced solvation of the mellitate species. Confirmation of the importance of electrostatic effects can be seen from the fact that the mellitic acid adsorption edges (Figure 3) intersect near the point of zero charge (pHPZC) for kaolinite, ca. 5.5. The surface is positively charged below the pHPZC but negatively charged above it; in each case, the magnitude of the

Adsorption of Mellitic Acid onto Kaolinite

charge increases with temperature, respectively enhancing or inhibiting the adsorption of negatively charged mellitate species. Similarly, Ward and Brady18 found that at higher temperatures the adsorption of oxalate to kaolinite increased below pH 6 but decreased above pH 6. Other studies of anion adsorption onto various oxides and soils28-31 have generally demonstrated a trend for adsorption to decrease with increasing temperature over a wide pH range. Since soils generally carry a net negative charge at pH values above 4, these results are also in accord with our study. ATR-FTIR Spectroscopy. The band pattern in the ATRFTIR spectra of adsorbed mellitic acid changes little from pH 7 to 4.9, with three bands at 1584, 1433, and 1338 cm-1. The spectrum at pH 3.9 is also similar, although the first and third bands are shifted somewhat to higher wavenumbers. These spectra have very similar band positions to the L6- solution spectrum, though all are shifted slightly to higher wavenumbers (ca. 6 cm-1). Thus, the spectra suggest that mellitic acid is adsorbed to kaolinite as the L6- ion at pH g 5. The relatively small shift in band positions and the similarity in band shapes to those in solution indicate that adsorption most probably involves outersphere complexation. Between pH 4 and 5, the further small shift in the positions of the outer bands most probably indicates the presence of HL5-. Not until the pH is near 3 do the surface spectra show evidence of the more-protonated H2L4- species, with the appearance of bands at 1712 and 1254 cm-1. Hence, the spectra are consistent with outer-sphere complexation of L6and HL5- over the pH range used in the adsorption experiments, as proposed in the surface complexation model. The increase in protonation of surface species compared with solution species at the same pH, especially at lower pH values, has been noted before for adsorption of various benzenecarboxylates onto goethite.6,9,27,32 While this may be the result of a higher pH adjacent to the positively charged surface, the fact that the surface species at pH 4 are less protonated on kaolinite than on goethite,27 although the positive charge on the kaolinite surface is smaller than that on goethite, suggests that surfacespecific factors are also important. While our infrared spectra indicate that outer-sphere complexation is the dominant mode of adsorption for mellitic acid on kaolinite, significant changes in adsorbed spectra led Specht and Frimmel11 to propose that oxalic, malonic, and succinic acids formed inner-sphere complexes with kaolinite. Their spectroscopic results are surprising, as they suggest significant adsorption of both oxalic and succinic acid above pH 10, while Ward and Brady18 found no significant adsorption of oxalate to kaolinite above pH 10. Surface Complexation Modeling. When modeling the adsorption of mellitic acid onto kaolinite, we sought to minimize the number of adjustable parameters while giving an adequate fit to the experimental data. We do not claim that the model presented here provides a complete description of all the surface processes that occur. In fact, the spectroscopic results indicate that it is almost certain that more highly protonated mellitic acid species would be involved in sorption below pH 3.5. Our aim was to find a set of surface reactions that adequately described the potentiometric titration and adsorption data at the five temperatures studied, while being consistent with the spectro(28) Barrow, N. J.; Shaw, T. C. Soil Sci. 1977, 124, 347-354. (29) Barrow, N. J. J. Soil Sci. 1986, 37, 267-275. (30) Barrow, N. J. J. Soil Sci. 1992, 43, 37-45. (31) Goldberg, S.; Forster, H. S.; Heick, E. L. Soil Sci. 1993, 156, 316-321. (32) Boily, J.-F.; Persson, P.; Sjo¨berg, S. J. Colloid Interface Sci. 2000, 227, 132-140. (33) CRC Handbook of Chemistry and Physics, 72nd ed.; Lide, D. R., Ed.; CRC Press: Boca Raton, FL, 1991.

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scopic data. The reactions and equilibrium constants generated by fitting the model to the data were then used to generate thermodynamic parameters which provide additional information about the adsorption processes. The proton stoichiometries measured during adsorption at constant pH support the reaction mechanisms proposed in the modeling. Speciation curves proposed by the model for mellitic acid adsorption on kaolinite are shown in Figure 4. We can estimate the proton stoichiometry predicted by the model for the kaolinite system at each pH from the surface and solution species present. The main surface reactions at pH 4.5 are

H2L4- + SOH h [(SOH2+)(LH5-)]4- (65%) H3L3- + SOH h [(SOH2+)(LH5-)]4- + H+ (35%) while, at pH 8.5, the reactions are

L6- + SOH h [(SOH)(L6-)]6- (50%) L6- + SO- + H+ h [(SOH)(L6-)]6- (50%) The predicted proton stoichiometries are therefore quite small (χ ≈ +0.3 at pH 4.5 and χ ≈ -0.5 at pH 8.5). The experimental values obtained from the isotherms show values of χ near zero at both pH values, consistent with the proposed reaction scheme. The dominant surface reaction at pH 8.5 suggests that mellitic acid is held to the kaolinite surface through hydrogen bonding to neutral SOH sites. This mechanism corresponds with that proposed by Ward and Brady18 for the adsorption of oxalic acid onto kaolinite. While at first sight this reaction may appear to contradict the importance of electrostatic interactions in determining the temperature dependence of adsorption, it need not. As the surface becomes more negatively charged at higher pH, there will be fewer SOH sites and the net negative charge will repel negatively charged mellitate ions, thereby decreasing adsorption. Previous models for the adsorption of benzenecarboxylates onto goethite have proposed outer-sphere complexation with SOH2+ sites only.6,9,19 For goethite, the relatively high pHPZC (∼8.5) means that significant concentrations of SOH2+ sites are available for adsorption even at relatively high pH. Thermodynamic Parameters. The thermodynamic parameters obtained from the surface complexation modeling of mellitic acid adsorption on kaolinite indicate that adsorption enthalpies are relatively small. On the other hand, the entropy change for the formation of [(SOH2+)(HL5-)]4-, estimated from the variation with temperature of the equilibrium constants, is relatively large. This most probably reflects a marked change in the hydration of both the mellitate ions and the charged surface upon adsorption. The reduction in the net charge of the complex compared with the separate reactants will result in a substantial liberation of water molecules. Not surprisingly, the entropy change for the formation of [(SOH)(L6-)]6- is much smaller, as there is no change in net charge on formation of that complex. It is difficult to compare the thermodynamic parameters determined in this study with those found in other work since the values depend both on the model used and, to some extent, on the experimental procedure employed. With these reservations, we make the following comments. In a study of organic acid adsorption onto kaolinite at two different temperatures, Ward and Brady18 deduced that the enthalpy change for the formation of an outer-sphere oxalate-kaolinite surface complex was small, in agreement with our results for the mellitate-kaolinite system. This small enthalpy change is consistent with the notion that

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mellitate-kaolinite bonds are the result of electrostatic and hydrogen-bonding interactions (outer-sphere complexation). The large entropy change found on formation of [(SOH2+)(LH5-)]4is consistent with the values found for the similar complexes formed on adsorption of mellitate onto goethite.19 This indicates that solvent effects are likely to play an important role in the adsorption of organic anions.

Conclusion Mellitic acid adsorption increases at higher temperature when the pH is below pHPZC, while at pH values above the pHPZC, adsorption decreases as temperature increases. This behavior most probably results from the effect of temperature on the surface charge of kaolinite: the net positive surface charge increases with increased temperature at low pH, but the net negative charge increases with increasing temperature at higher pH. Thus, the relative concentration of surface sites on which adsorption occurs is the major determinant of the adsorption of mellitic acid onto kaolinite. ATR-FTIR spectroscopy and surface complexation modeling suggest that mellitic acid is bound to kaolinite as outer-sphere complexes in the pH range from 3 to 8. The principal adsorbed species at lower pH, [(SOH2+)(HL5-)]4-, occurs over a range of pH where solution species are protonated to a much greater

AngoVe et al.

extent. While electrostatic attraction is involved in sorption at lower pH, the species proposed at higher pH, [(SOH)(L6-)]6-, involves hydrogen bonding between the mellitate ion and neutral surface species. The enthalpies of adsorption of both species are small, but the entropy of adsorption of the lower-pH species is large and positive, suggesting that solvent effects play an important role. Not surprisingly, the permanent negative-charge sites on kaolinite play no role in sorption. These results indicate that polar organic species adsorb to SOH groups on clay minerals principally below pH 8. Thus, at higher pH values, both surface hydroxyl and permanent-charge sites are available to interact with metal cations. Since polar organic species bind by outer-sphere complexation, their carboxylate groups are largely free to complex with metal ions. Sorbed humic and fulvic acids are therefore expected to increase metal ion adsorption at lower pH, while causing little change to sorption at higher pH values. Acknowledgment. Financial support was provided by the Australian Research Council Small Grants Scheme. Professor Staffan Sjo¨berg is gratefully acknowledged for his contribution and suggestions. LA0534571