Desorption of Humic Acids from an Iron Oxide Surface

Wageningen Agricultural University, P.O. Box 8038,. 6700 EK Wageningen, The Netherlands. The adsorption and desorption of different humic acids on ...
14 downloads 0 Views 84KB Size
Download PDF
Environ. Sci. Technol. 1998, 32, 2572-2577

Desorption of Humic Acids from an Iron Oxide Surface M A R C E L O J . A V E N A * ,† A N D LUUK K. KOOPAL Laboratory for Physical Chemistry & Colloid Science, Wageningen Agricultural University, P.O. Box 8038, 6700 EK Wageningen, The Netherlands

The adsorption and desorption of different humic acids on iron oxide-coated silicon plates was studied in stagnation point flow using a reflectometric technique. The adsorption increased by decreasing the pH. Desorption studies were conducted by first adsorbing the humics at low pH and then desorbing them by either increasing the pH or rinsing with supporting electrolyte (0.1 M KNO3) at a constant pH. Two different desorption processes can be distinguished depending on the method used for desorbing the humics: desorption upon change in pH without changing the HA concentration and desorption upon dilution without changing the pH. When desorption is induced by changing the pH, the desorption is fast and reversible, and the rate of desorption increases by increasing the pH. At the new pH the surface becomes oversaturated with HA molecules because of increased molecule-molecule repulsion (electrostatic) and decreased surface-molecule attraction (electrostatic and specific). Desorption upon dilution at a constant pH is so slow that the equilibrium condition cannot be reached in practice. Since there is no change in pH, the interactions in the adsorbed layer remain unchanged, and the only driving force for desorption is given by the decrease of the HA concentration in the bulk solution. Such a change is much less effective than a direct change of the interactions in the adsorbed layer (by changing the pH), and it is insufficient to promote a fast desorption.

Introduction The interaction between the polymeric soil organic matter and oxide surfaces is a widely studied subject because of its relevance to many aspects of soil chemistry and environmental chemistry. Humic acids (HA) and fulvic acids (FA) are very active in binding ions and other organic molecules, and they have a relatively strong affinity for oxide surfaces. These properties make humic and fulvic acids very important in regulating the speciation and mobility of ions, nutrients, and contaminants in soils and aquifers (1, 2). The adsorption of HA on different oxide or clay surfaces has been studied as a function of pH, ionic strength, HA concentration, type of humic, etc. (3-10). The adsorption is governed by specific interactions between HA reactive groups (mainly carboxylic and phenolic) and the reactive groups of the oxide surfaces and by electrostatic interactions (5, 8, 9). The combination of both effects generally leads to an increase in the adsorbed amount by decreasing the pH. * Corresponding author phone: (31) 317 482482; fax: (31) 317 483777; e-mail: [email protected]. † On leave of absence from the Departamento de Fisicoquı ´mica, Facultad de Ciencias Quı´micas, Universidad Nacional de Co´rdoba, Co´rdoba, Argentina. 2572

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 32, NO. 17, 1998

Desorption of HA from solid surfaces is less well studied (5-8, 10), even though the study of desorption phenomena is as important as adsorption studies for the understanding of the mechanisms that govern the interaction between HA and oxides. Gu et al. (6, 8) and Zhou et al. (7) studied desorption upon dilution (by decreasing the HA concentration or washing with supporting electrolyte but without changing the pH of the dispersion) whereas Davis (5) and Vermeer and Koopal (10) studied desorption upon changes in pH (by changing the pH of the dispersion to a less favorable condition for adsorption) without modifying the HA concentration. The main conclusion that emerges from studies of desorption upon dilution is that, once adsorbed at a given pH, the soil organic matter is very difficult to desorb by dilution at that pH. The desorption and adsorption isotherms of a given HA do not coincide: the desorption curve lies considerably above the adsorption one and resembles a highaffinity isotherm as compared with the generally observed “rounded” adsorption isotherm that continues to increase at increasing HA concentration (7, 8, 10). This deviation between the adsorption and desorption curves is called hysteresis and exists even after equilibration for extended times, for instance, as long as 63 days (8). Several explanations have been offered for the hysteresis. Vaccari and Kaouris (11) have attributed this phenomena to surface heterogeneity, assuming that two types of sites are available on the sorbent; sites on which adsorption is irreversible and sites where adsorption is reversible. Zhou et al. (7), on the other hand, explained the difference in the desorption capacity and sorption hysteresis of several humics substances assuming differences in the interaction mechanism; some humic molecules should be adsorbed chemically and irreversibly, while other humic molecules were supposed to be adsorbed physically and reversibly by electrostatic forces. Gu et al. (6, 8) proposed two explanations: (1) the existence of multiple binding between adsorbed molecules and the iron oxide surface, so that desorption requires simultaneous detachment of all bound segments, and (2) the existence of heterogeneity of the adsorbate having components with different affinity for the surface, so that surface fractionation takes place (the composition of the adsorbed layer results to be different from that of the solution) and the surface becomes enriched with high affinity components. Because of this surface fractionation, the composition of the solution during adsorption will be different from that at desorption leading to a difference between the adsorption and the desorption isotherms (8). This last theory of surface fractionation is similar to the polydispersity theory used to explain desorption hysteresis of polymers (12, 13). This theory was extended for polyelectrolytes by Vermeer and Koopal (10) and used to explain the hysteresis of HA adsorption. According to this theory, the differences in affinity are mainly due to polydispersity with respect to molecular weight, and it is not necessary to address chemical heterogeneity of the sample. The theory predicts that, at a given pH and salt concentration, polyelectrolytes of high molecular weight have a higher affinity for the surface than low molecular weight ones because big molecules lose less translational entropy when adsorbing (10, 12). When a polydisperse polyelectrolyte adsorbs, surface fractionation takes place since relatively small molecules adsorb only when the amount of big molecules is not enough to saturate the surface (i.e., at low overall concentration of polyelectrolyte). When the supply of big molecules is high, the relatively fast adsorbing lower molecular weight fraction is displaced at longer adsorption S0013-936X(98)00112-6 CCC: $15.00

 1998 American Chemical Society Published on Web 07/25/1998

times by the big molecules. At high concentrations, therefore, displacement results in an adsorbed layer that is almost exclusively composed of the high molecular weight fraction. Upon dilution, the desorption isotherm follows the highaffinity isotherm of the high molecular weight species rather than the “rounded” isotherm of the mixture, and this produces hysteresis. In this situation, the hysteresis is not a nonequilibrium phenomenon (10, 12). For polyelectrolytes, an alternative to desorption upon dilution is desorption upon changes in pH at a constant HA concentration. The main conclusion that emerges from studies of Davis (5) and Vermeer and Koopal (10) is that the organic matter desorbs rapidly from the surface when the pH is changed to a less favorable condition. The desorption curve approximates the adsorption one after 1 h of equilibration, and no large hysteresis is found. The experimental evidence for this type of desorption is, however, limited. The fact that the hysteresis of desorption upon dilution is quite different from that at desorption by changing the pH indicates that quite different mechanisms for desorption are involved, and both processes need to be considered in more detail. The aim of this paper is to analyze the adsorption and desorption behavior of three different HA samples on an iron oxide surface, with emphasis to the desorption aspects. The amount ad- or desorbed is followed as a function of time and pH at a given humic acid and electrolyte concentration. The desorption rate upon dilution with pure solvent at a constant pH is also studied. In both cases the behavior is analyzed on the basis of modern insights in polyelectrolyte adsorption.

of the Fe2O3 layer was around 5 as revealed by streaming potential measurements and by adsorption of positively and negatively charged strong polyelectrolytes. The iep value is lower than the iep or point of zero charge of crystalline iron oxides (18). However, this behavior is not exclusive for iron oxide coatings. Relatively low iep values have also been reported for titanium and iridium coatings on silica (19, 20). According to Giatti and Koopal (19), this could be due to a surface that is less well structured than that of ordinary crystalline oxides. Reflectometry. Adsorption-desorption measurements were performed by reflectometry, in a stagnation point flow cell as described by Dijt et al. (21). The principle of reflectometry is as follows. A laser emits a polarized beam, which is reflected by the wafer in the reflectometric cell. The reflected light is split into its parallel and perpendicular components, which are detected separately with photodiodes. The signal S is defined as

S)f

(1)

where f is an equipment constant and Rp and Rs are the intensities of the parallel and the perpendicular components, respectively. Upon adsorption of a polymer on the oxide layer, the signal changes by an amount ∆S. In adequate conditions, the relative change of the signal, ∆S/S0, is proportional to the adsorbed amount (22)

Γ)

Materials and Methods Chemicals were p.a. quality, and high purity water was used. All experiments were performed at constant temperature, 22 °C. Humic Acids. Three different humic acids were used: purified Aldrich humic acid (PAHA), Shitara black humic acid (SBHA), and Kinshozan P humic acid (KPHA). The origins, methods of extraction, and purification of these samples together with some physical and chemical properties were reported respectively in previous papers (14-16). The elemental analysis of PAHA is (wt %) C, 55.8%; O, 38.9%; H, 4.6%; and N, 0.6%. That of SBHA is C, 58.7%; O, 34.5%; H, 3.4%; and N, 3.4%. That of KPHA is C, 57.3%; O, 36.3%; H, 3.8% and N, 2.6%. The total acidities of PAHA, SBHA, and KPHA are 5, 6, and 5.5 mmol/g, respectively. In a previous study by viscometry, it was shown that the studied humics behave as flexible entities that can swell or shrink in response to changes in pH and ionic strength (17). Increasing the ionic strength or decreasing the pH lead to molecular shrinkage. However, a comparison with the behavior of branched and linear polyelectrolytes indicates that the swelling and shrinking capacity of the humics is small, similar to that of branched molecules, because of their internal structure that limits expansion-contraction processes (17). Before use, a stock solution (2-3 g/L) was prepared by dissolving the samples in KOH at pH between 10 and 11. The working solutions (usually 50 mg/L) were made by diluting the stock solution with the supporting electrolyte, 0.1 M KNO3, and the pH was adjusted to the desired value with KOH and HNO3 solutions. Iron Oxide. Strips of silicon wafers bearing a layer of Fe2O3 were used as adsorbent. The wafers, which are of the Czochralsky type, were obtained from Aurel GmbH (Germany). The iron oxide layer was deposited by reactive sputtering of Fe in an oxygen atmosphere. This was carried out at Phillips Laboratories at Eindhoven, The Netherlands. Electron microscopy and atomic force microscopy revealed that the surface of the Fe2O3 layer is very smooth and flat, with no cracks or irregularities. The isoelectric point (iep)

Rp Rs

1 ∆S As S0

(2)

where As is the sensitivity factor. This factor was calculated using the method of Hansen (23), which is based in the matrix formalism of Abeles. As depends on the thickness d of the oxide layers and on the refractive indexes n of silicon, the oxide layers, the adsorbed layer, and the solution. Also the refractive index increment dn/dc of the adsorbed layer (c stands for HA concentration in g/cm3), the angle of incidence θi, and the wavelength λ of the laser beam must be known. The values used were nSi ) 3.8, nSiO2 ) 1.46, nFe2O3 ) 2.966, nHA ) 1.363, nwater ) 1.333, dSiO2 ) 80 nm, dFe2O3 ) 55 nm, dn/dc ) 0.28 cm3/g, θi ) 70.6, λ ) 632.8 nm. The refractive index increment was determined with a differential refractometer; d and n values were determined by ellipsometry. For the present set of parameters, As is -0.045 m2/mg. A stagnation point flow is obtained by flowing the solvent or the HA solutions (about 1 mL/min) into the cell through a cylindrical channel perpendicular to the wafer. The intersection of the symmetry axis of the cylinder with the surface is the stagnation point (21). This point is positioned such that it coincides with the reflection spot of the laser beam. With the help of a two-way valve, it is possible to switch from solvent to HA solution or vice versa. For more information about reflectometry and stagnation point flow, the study of Dijt et al. can be consulted (21).

Results and Discussion Figure 1 shows an example of a typical adsorption-desorption experiment due to a change in pH. Initially, only supporting electrolyte (pH ) 3.25, 0.1 M KNO3) is flowing, and a stable baseline is obtained. The noise and the detection limit (approximately 0.05 mg/m2) are low. Once the flow is switched to a 50 mg/L HA solution (same pH and electrolyte concentration), the signal changes strongly because adsorption takes place. The adsorption occurs rapidly in the first 100-200 s. After that, there is a decrease in the adsorption rate and it increases slowly with time over a much large time range (up to several hours). The fast process is due to direct adsorption, whereas the slow process may be related to VOL. 32, NO. 17, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

2573

FIGURE 1. Adsorption-desorption experiment for PAHA. Adsorbing solution: 50 mg/L HA, 0.1 M KNO3, pH ) 3.25. Desorbing solution: 50 mg/L HA, 0.1 M KNO3, pH ) 7.8.

FIGURE 2. Desorption kinetics of PAHA at different pH. Adsorbing solution: 50 mg/L HA, 0.1 M KNO3, pH ) 3.25. Adsorption time: 1000-1500 s. Desorbing solution: 50 mg/L HA, 0.1 M KNO3, pH is indicated in the figure. rearrangements in the adsorbed layer either both by conformational changes in the adsorbed layer and/or by exchange of adsorbed molecules with newly arriving molecules. Although the kinetics of HA adsorption is a very important subject, it is not the main topic of this paper, and it will be considered in another publication together with equilibrium adsorption data. After 1500 s of adsorption, the flow was switched to a HA solution of the same humic concentration and ionic strength but a different pH (7.8), where the adsorption is less favorable. The sharp increase in the amount adsorbed is only an artifact that is sometimes noted when the valve is switched to flow a different solution. A relative fast desorption can be observed, and after about 300 s the signal seems to stabilize at a constant value. Although part of the humic remains adsorbed, the data in Figure 1 indicate that an important fraction can be desorbed readily by changing the pH. The detailed effects of the desorption pH can be seen in Figure 2, which shows the desorption curves for PAHA initially adsorbed at pH ) 3.25 (0.1 M KNO3) and desorbed by flowing humic acid solutions of the same HA and salt concentrations but with pH values ranging from 4.8 to 10.75. Except for pH 4.8, the desorption process is fairly fast and shows stabilization to a constant value. At pH ) 4.8, the desorption is relatively slow, and there is no stabilization yet after 1000 s. Comparison of the initial slopes shows that the desorption rate is clearly pH dependent, increasing as the pH increases. The behavior under the same conditions but in absence of humics in the desorbing solution is shown in Figure 3. A similar trend can be observed, although the desorption rate is substantially lower and for pH < 6 the system needed more than 1000 s to reach the final desorption level. 2574

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 32, NO. 17, 1998

FIGURE 3. Desorption kinetics of PAHA at different pH. Adsorbing solution: 50 mg/L HA, 0.1 M KNO3, pH ) 3.25. Adsorption time: 1000-1500 s. Desorbing solution: 0.1 M KNO3, pH is indicated in the figure. Conclusions about the adsorption level for times greater than about 1000 s become less reliable because a gradual baseline drift that is comparable to the slope of the measurement curves might occur. For the moment no clear explanation can be provided for the lower desorption rates observed in Figure 3 as compared to those in Figure 2. A reason might be that some incoming HA molecules temporarily adsorb with a few segments to the surface, thereby increasing the negative charge in the adsorbed layer and promoting the other molecules to desorb faster. To check reversibility, Figure 4 compares the adsorptiondesorption levels for the three HA samples for desorption upon changes in pH. The adsorbed amount at different pH values was obtained by flowing the HA solution at the desired pH during 1000-1500 s (solid symbols). The desorption levels were obtained by first adsorbing HA in electrolyte at pH 3.254.00 and then desorbing the humics with either a HAelectrolyte solution (open squares) or a pure electrolyte solution (open triangles) both at the desired pH. The total desorption time was 1000 s in all cases, and the KNO3 concentration was 0.1 M. The adsorbed amount that remains after the desorption treatment reaches values that are very close to those obtained by adsorbing the humics at the pH under investigation. A closer look at the data in Figure 4 indicates that in some cases, Γ after desorption is slightly higher than the corresponding Γ obtained by adsorption. However, it is expected that both values would converge to the same level if somewhat more time for desorption had been allowed. For the explanation of the adsorption as a function of pH, we refer to the work of Vermeer et al. (9, 10) where the HA acid adsorption behavior is compared with theoretical predictions of weak polyelectrolytes adsorption on a variable charge surface. As has been outlined (9, 10), the Γ vs pH behavior must be governed by both electrostatic and specific interactions. Electrostatic interactions must be present because both the surface and the particles carry electrical charges. The adsorption will be favored by electrostatics at pH < iep of the surface and disfavored at pH > iep. Specific interactions must also be involved since otherwise HA would not adsorb at pH > iep. Specific interactions are usually considered to be complexation reactions of carboxilic and phenolic groups of the HA with the surface hydroxyl groups of the oxide. The fact that adsorption and desorption curves are similar when HA is present in the desorbing solution suggests that the processes involved in these conditions are fast and reversible. The non-desorbed fraction observed in Figures 1, 2, and 4 is the normal adsorbed amount at the pH and ionic strength of the desorption experiment. This behavior is similar to that found by Davis (5).

FIGURE 5. Schematic representation of the desorption processes.

FIGURE 6. PAHA. Desorption kinetics upon dilution with pure solvent. Adsorbing solution: 50 mg/L HA, 0.1 M KNO3, pH ) 4.0. Desorbing solution: 0.1 M KNO3, pH ) 4.0. Adsorption times are indicated in the figure.

FIGURE 4. Amount of adsorbed HA after adsorption (solid symbols) and desorption (open symbols) experiments for the three HA studied. 0.1 M KNO3: triangles, desorption in absence of HA; squares, desorption in the presence of HA. Desorption time: 1000 s. Adsorption pH: PAHA, 3.25; SBHA, 3.50; KPHA, 4.00. The error bars represent the standard deviation obtained by repeating 20 times the adsorption experiment at pH ) 4. When the desorption is done in the absence of HA, the adsorbed amount should drop to zero because there is no HA present in the flowing solution and every desorbed molecule is removed by the flux. However, desorption takes place quickly only until Γ reaches the value that would corresponds to a HA solution in equilibrium with the surface at the desorption pH. In other words, once Γ decreased to a value close to that described by the curve in Figure 4, the desorption becomes very slow. The particular desorption behavior observed for the humics studied suggests that two desorption processes are involved. The two processes are schematically represented in Figure 5, where the measured Γ vs pH curve divides the desorption domain in two regions. If the system is found in the region above the curve, like the point denoted as 1 (obtained by changing the pH from low to high values, horizontal arrow), a rapid desorption takes

place until the Γ vs pH curve is reached (point 2). In the presence of HA in solution this is the final situation. In the absence of humics in the desorbing solution the system would tend to equilibrate by reaching point 3, but the desorption takes place at a very slow rate. Therefore, in practice desorption upon a change in pH leads to adsorption values that coincide with the adsorption curve, even when the desorption is performed with pure electrolyte. According to the discussion presented above, desorption upon dilution without a pH change is expected to be very slow (e.g., from point 2 to point 3). This was tested with experiments in which the desorbing electrolyte solution had the same pH as the adsorbing solution. The time allowed for desorption was around 6000 s. Figure 6 shows in a semilogarithmic plot the adsorbed fraction as a function of the desorption time for PAHA samples that were allowed to adsorb at pH 4 for 200 s, 2000 s, and 16 h prior to the desorption experiment. The data indicate very slow desorption in all cases, but the longer the adsorption times, the slower the desorption becomes. If the last part of the plots (t > 200 s) is assumed to be linear and extrapolated, it can be estimated that around 10 days would be necessary to desorb 35% of the fastest desorbing sample (adsorption time 200 s). A 50% desorption of this sample could be reached after 1.3 years. Similarly a desorption of 40% and 18% of the samples adsorbed during 2000 s and 16 h, respectively, is predicted in 1.3 years. Clearly, the desorption rate is so low that the equilibrium condition (Γ ) 0) cannot be reached in practice. Although not shown here, very slow desorption upon dilution was also found for KPHA and SBHA. This behavior is in agreement with data shown by Gu et al. (6, 8) and Zhou et al. (7), and it indicates that it is very difficult to desorb HA by dilution at a constant pH. The variation in the desorption rate by varying the adsorption time from 200 s VOL. 32, NO. 17, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

2575

to 16 h is evidence of either or both rearrangement in the adsorbed layer by conformational changes and/or progressive surface fractionation. In both cases, the molecules will be more strongly bound to the surface at longer times, and this makes the desorption more difficult. In practice, where often the humic acids have been in contact with the mineral particles for a long time, it can be concluded that desorption upon dilution can be neglected for time scales of months. As mentioned before, the fact that desorption upon a change in pH and desorption upon dilution have rather different time scales points toward different desorption mechanisms. The desorption mechanism that takes place upon an increase in the pH is related to a change in the interaction energies between the HA molecules and the surface and among the HA molecules laterally. An increase in pH involves (1) a decrease of the positive charge of the surface or detaching of carboxylic and phenolic groups from the surface groups through a fast exchange with hydroxyl ions and (2) an increase of the negative charge of the humic molecules that leads to mutual repulsion. Together the two changes lead to a net repulsion and to a fast moving of the HA molecules away from the surface. One could also say that at the new pH the surface becomes oversaturated with HA molecules because of increased molecule-molecule repulsion (electrostatic) and decreased surface-molecule attraction (electrostatic and specific). When HA molecules are present in the desorbing solution, the desorption takes place until the system reaches the equilibrium adsorbed state corresponding to the new pH (point 2 in Figure 5), where a new balance of the electrostatic and specific interactions is found. This desorption mechanism explains the pH dependence of the desorption rate: the higher the pH, the more effective are the two effects that promote desorption. When HA molecules are not present, the system also reaches point 2 (see Figure 5) quickly because the same repulsion forces are exerted in the adsorbed layer. Moving from point 2 to point 3 can be seen as desorption upon dilution at a constant pH. The desorption mechanism for desorption upon dilution at a constant pH is totally different than desorption by a pH change. Since there is no change in pH, the interactions in the adsorbed layer remain unchanged, there is no surplus of hydroxyls to exchange with carboxylic and phenolic groups, and there is no increase in the negative charge of the humics to stimulate the moving of the molecules away from the surface. The only driving force for desorption is given by the decrease of the HA concentration in the bulk solution. Obviously such a change is much less effective than a direct change of the interactions in the adsorbed layer, and it is insufficient to promote a fast desorption. A quantitative explanation for the low rate of desorption upon dilution of uncharged polymers has been suggested by Dijt et al. (13). In this model it is assumed that attachmentdetachment reactions at the surface are fast and that adsorbed molecules are in equilibrium with those in a subsurface layer (adjacent to the surface). Thus, the surface is continuously in a local equilibrium with the molecules in the subsurface, and the desorption or adsorption kinetics are given by the rate of transport of molecules from the subsurface layer to the bulk or vice versa. Under these conditions, the adsorption or desorption rate is simply given by

dΓ ) cb - ceq(Γ) dt

R

(3)

where R is an adsorption barrier, cb is the bulk concentration of polymer, and ceq(Γ) is the equilibrium concentration in the subsurface layer corresponding to a particular value of Γ. For adsorption cb > ceq and for desorption cb < ceq. The function ceq(Γ) is the inverse of the (high affinity) adsorption 2576

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 32, NO. 17, 1998

FIGURE 7. Schematic illustration of desorption rates as deduced by using the local equilibrium concept. The full curve is the inverse of an hypothetical equilibrium adsorption isotherm since the desorption rate is proportional to ceq. isotherm of the adsorbed polymer. For uncharged polymers, R depends on the hydrodynamic conditions and on the diffusion coefficient of the adsorbate (21). For polyelectrolytes, eq 3 only applies at a given pH and salt concentration, and R includes an electrostatic term that depends on the adsorption itself (24). In the particular case of desorption upon dilution by flowing a solution with cb ) 0, RdΓ/dt equals -ceq(Γ), the negative sign indicating desorption. Therefore, at a fixed flow rate and negligible variations in R, the desorption rate is only a function of the adsorbed amount at equilibrium and has the shape of an inverted isotherm, as shown schematically in Figure 7. The shape of the inverted isotherm corresponds to the isotherm of the adsorbed species, which for a polydisperse sample is the high affinity isotherm of the preferentially bound molecules. When the system is near saturation (somewhere in the vertical part of the curve), the desorption rate is high. However, once a small fraction has been desorbed, the subsurface concentration ceq drops strongly; hence, the rate decays considerably. If Γ has decreased so that the horizontal part of the curve is reached, the desorption rate becomes unmeasurably slow. It should be noted that in the case of desorption the assumption that the rate of detachment of the molecules is fast may not be correct. However, a slower detachment rate only slows down the desorption even further. Based on eq 3 and theoretical prediction of ceq(Γ) Dijt et al. (13) derived that the amount of molecules that remain adsorbed must decrease approximately logarithmically with t. On the basis of this approximation, the extrapolated desorption results after 10 days and 1.3 years (see Figure 6) have been obtained. The previous results and discussion suggest that the proposal of Vaccari and Kaouris (11) is not valid to explain the desorption behavior of HA on iron oxides. If the hysteresis were only due to the presence of sites where the adsorption is reversible and sites where the adsorption is irreversible, the desorption curve in Figure 4 would not be a line that corresponds to the adsorption curve, but a horizontal line corresponding with the irreversible adsorbed amount. The same analysis allows us to discard the theory of Zhou et al. (7), which also assumes the presence of molecules that adsorb reversibly and molecules that adsorb irreversibly. The multiple binding theory proposed by Gu et al. (8) is correct in the sense that the high affinity of the HA for the surface that lead to multiple binding makes the desorption upon dilution difficult. However, the theory does not address the desorption by a pH change. In fact, if the detachment of all surface-bound segments was practically impossible under all conditions, the desorption rate in the region above the Γ vs pH curve would be also very slow, and the coincidence

between the adsorption and desorption curves would not occur. The surface fractionation or polydispersity theory was shown to be adequate to explain adsorption hysteresis upon dilution in the case of polymers (12) and polyelectrolytes (10). Since mostly molecules with very high affinity for the surface are adsorbed, a small desorption of these molecules is enough to increase the concentration in the aqueous solution and to create a new equilibrium situation that stops desorption. However, this theory is only valid when batch experiments are performed, where the desorbed molecules contribute to increase the concentration in solution. In a flux system such as the one used here, every desorbed molecule is removed by the flux, and the only equilibrium condition is Γ ) 0 when rinsing with electrolyte. Thus, a model that predicts low desorption kinetics such as the one schematized by Figure 7 needs to be considered. In this last case, the polydispersity effect is still relevant because it explains why the adsorption isotherm that governs the desorption rate is of the high affinity type. In general, the present study on the desorption behavior of humics on iron oxide indicates that in practice desorption cannot be achieved successfully unless a change in pH has occurred. For example, a HA that is adsorbed on an oxide at a relatively low pH can be partially desorbed or mobilized by an increase in pH. Similarly as discussed for the pH, a change in the salt concentration (at a constant pH) could also lead to some desorption. The magnitude of this effect will depend on how the change in the electrolyte concentration changes the interactions in the adsorbed layer.

Acknowledgments M.J.A. thanks the Wageningen Agricultural University and the Laboratory for Physical Chemistry & Colloid Science (WAU) for the first year postdoctoral fellowship and CONICET for the second one. This work was carried out as part of the European Union Project Intas-Ukraine 95-0165.

Literature Cited (1) McCarthy, J. F.; Zachara, J. M. Environ. Sci. Technol. 1989, 23, 496-502. (2) Murphy, E. M.; Zachara, J. M.; Smith, S. C. Environ. Sci. Technol. 1990, 24, 1507-1516. (3) Davis, J. A. Geochim. Cosmochim. Acta 1982, 46, 2381-2393.

(4) Varadachari, C.; Chattopadhyay, T.; Ghosh, K. Soil Sci. 1997, 162, 28-34. (5) Davis, J. A. In Contaminants and Sediments; Baker R. A., Ed.; Ann Arbor Science: Ann Arbor, 1981; Vol. 2, Chapter 15. (6) Gu, B.; Schmitt, J.; Chen, Z.; Llang, L.; McCarthy, J. F. Geochim. Cosmochim. Acta 1995, 59, 219-229. (7) Zhou, J. L.; Rowland, S. J.; Fauzi, R.; Mantoura, R. F. C.; Braven, J. Water Res. 1994, 28, 571-579. (8) Gu, B.; Schmitt, J.; Chen, Z.; Llang, L.; McCarthy, J. F. Environ. Sci. Technol. 1994, 28, 38-46. (9) Vermeer, A. W. P.; van Riemsdijk, W. H.; Koopal, L. K. Langmuir 1998, 14, 2810-2819. (10) Vermeer, A. W. P.; Koopal, L. K. Langmuir In press. (11) Vaccari, D. A.; Kaouris, M. J. Environ. Sci. Health 1988, A23, 797-822. (12) Fleer, G. J.; Cohen Stuart, M. A.; Scheutjens, J. M. H. M.; Cosgrove, T.; Vincent, V. Polymers at Interfaces; Chapman & Hall: London, 1993. (13) Dijt, J. C.; Cohen Stuart, M. A.; Fleer, G. J. In Colloid-Polymer Interactions Dublin, P. L., Tong, P., Eds.; ACS Symposium Series No. 532; American Chemical Society: Washington, DC, 1993; Chapter 2. (14) Vermeer, A. W. P. Ph.D. Thesis, Agricultural University of Wageningen, 1996. (15) Kuwatzuka, S.; Tsutsuki, K.; Kumada, K. Soil Sci. Plant Nutr. 1978, 24, 337-347. (16) Tsutsuki, K.; Kuwatzuka, S. Soil Sci. Plant Nutr. 1978, 24, 547560. (17) Avena, M. J.; Vermeer, A. W. P.; Koopal, L. K. Colloids Surf. In press. (18) Lyklema, J. Fundamentals of Interface and Colloid Science, Volume II; Academic Press: London, 1995; Appendix 3. (19) Giatti, A.; Koopal, L. K. J. Electroanal. Chem. 1993, 352, 107118. (20) Hoogeveen, N. G.; Cohen Stuart, M. A.; Fleer, G. J. J. Colloid Interface Sci. 1996, 182, 133-145. (21) Dijt, J. C.; Cohen Stuart, M. A.; Hofman, J. E.; Fleer, G. J. Colloids Surf. 51, 141-158. (22) Dijt, J. C. Ph.D. Thesis, Agricultural University of Wageningen, 1993. (23) Hansen, W. N. J. Opt. Soc. Am. 1968, 58, 380. (24) Cohen Stuart, M. A.; Hoogendam, C. W.; de Keiser, A. J. Phys.: Condens. Matter 1997, 9, 7767-7783.

Received for review February 4, 1998. Revised manuscript received June 3, 1998. Accepted June 10, 1998. ES980112E

VOL. 32, NO. 17, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

2577