Langmuir 2005, 21, 11863-11869
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Separating Surface and Solvent Effects and the Notion of Critical Adsorption Energy in the Adsorption of Phenolic Compounds by Activated Carbons P. J. M. Carrott,* P. A. M. Moura˜o, M. M. L. Ribeiro Carrott, and E. M. Gonc¸ alves Centro de Quı´mica de E Ä vora and Departamento de Quı´mica, Universidade de E Ä vora, Cole´ gio Luı´s Anto´ nio Verney, 7000-671 E Ä vora, Portugal Received July 29, 2005. In Final Form: September 14, 2005 A modified form of the Freundlich equation in which the solute equilibrium concentration is normalized with respect to the solute solubility is analyzed and applied to adsorption isotherms of phenol, 4-nitrophenol, 4-chlorophenol, and 2-chlorophenol at different values of pH on commercial activated carbon before and after oxidation. The analysis confirms the importance of normalizing the solute equilibrium concentration when analyzing the adsorption isotherms, and it is suggested that a parameter, KF10, obtained by taking 10% solubility as the reference point when applying the Freundlich equation, is probably the best comparative estimate of the relative adsorption capacity of the carbon for different phenolic compounds. In combination with the Freundlich exponent, nF, estimates of the adsorption capacity at any other reference point can then be obtained. Analysis of the experimental results also indicates a need to distinguish between two regimes of adsorption, characterized by an adsorption energy, Eads, greater than or less than a critical value, Eca. When Eads > Eca, the shape of the adsorption isotherm is determined by solute-solid interactions. On the other hand, when Eads < Eca, solute-solution interactions become more important.
Introduction The adsorption of dissolved organic compounds by solid adsorbents is a complex process involving electrostatic, dispersive, and chemical interactions and is influenced by a large number of factors which include intrinsic properties of the solute (such as solubility and ionization constant), intrinsic properties of the adsorbent (such as point of zero charge and pore size distribution), and solution properties (in particular, pH) as well as by the temperature of the system.1 Excellent reviews of the literature pertaining to adsorption by carbon adsorbents were published relatively recently2,3 and highlighted the fundamental importance of adsorbent surface chemistry. In adsorption science in general, a commonly used method for quantifying the influence of surface chemistry on adsorption properties is by way of the analysis of adsorption isotherms using appropriate theoretical or empirical adsorption isotherm equations. Hopefully, variations in the derived parameter values can then be related to alterations in the surface chemistry of the adsorbent or in other properties of the system. Perhaps the most commonly used procedure for analyzing solute adsorption isotherms involves the application of the Freundlich equation which, for the present purposes, is most conveniently expressed in the form
nads ) KFCeq1/nF
(1)
where nads is the measured adsorption at a solute equilibrium concentration Ceq and the constant KF and * To whom all correspondence should be addressed. E-mail:
[email protected]. (1) Moreno-Castilla, C. Carbon 2004, 42, 83. (2) Radovic, L. R.; Moreno-Castilla, C.; Rivera-Utrilla, J. In Chemistry and Physics of Carbon; Radovic, L. R., Ed.; Marcel Dekker: New York, 2000; Vol. 27, p 227. (3) Dabrowski, A.; Podkoscielny, P.; Hubicki, Z.; Barczak, M. Chemosphere 2005, 58, 1049.
exponent nF are adjustable parameters whose values vary from one system to another. Equation 1 has been found to provide a very good fit, over a range of concentrations spanning several orders of magnitude, to adsorption isotherms given by a wide range of solutes on different adsorbents and, in the case of carbon adsorbents, for example, it is the basis for the standard method for evaluating the adsorption performance of industrial activated carbons.4 Despite its wide applicability, eq 1 suffers from a number of drawbacks. Perhaps the most significant of these are that the value of the constant KF depends on the units used to express Ceq (as well as nads), the equation does not give an estimate of the effective overall adsorption capacity of the adsorbent, and it is not normalized with respect to a standard state of the solute. The combination of these factors can sometimes make it difficult to compare results reported by different authors and also raises a doubt in relation to the precise significance of the parameters KF and nF. The main objective of this paper is to consider a methodology, based on a slightly modified form of eq 1, which circumvents these difficulties. Experimental Section 1. Materials. For this work, we studied the adsorption at 298 K and at different values of pH of a number of selected monophenolic compounds on commercial activated carbon before and after oxidation. The activated carbon was NORIT SX+ supplied by the manufacturers in a 1 kg polypropylene container. A portion of the activated carbon was oxidized by heating in nitric acid at 423 K for 3 h. It was then washed with distilled water until constant pH and dried at 423 K for 2 h. The as-received and oxidized samples are designated by NSXar and NSXox, respectively. Nitrogen adsorption at 77 K, after outgassing the samples at 573 K, was carried out using a CE Instruments Sorptomatic 1990. Analysis of the adsorption isotherms was carried out by means of the BET and DR methods.5,6 Elemental analysis of C, (4) ASTM Standard D3860-98 (Reapproved 2003).
10.1021/la0520886 CCC: $30.25 © 2005 American Chemical Society Published on Web 10/20/2005
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Table 1. Elemental, Surface Chemical and Textural Analysis of the Carbon Samples units elemental analysis C H N S O difference surface chemical analysis pzc acid sites basic sites textural analysis ABET vo (DR method) Lo (DR method)
% % % % % %
NSXar
NSXox
76.1 0.4 0.4 0.0 15.2 7.9
62.2 0.8 1.1 0.0 31.5 4.4
pH meqg-1 meqg-1
7.2 0.45 0.26
4.1 1.48 0.25
m2g-1 cm3g-1 nm
814 0.32 1.01
777 0.29 1.18
Figure 1. Adsorption isotherms (in any units of Ceq and nads) simulated on the basis of the Freundlich equation assuming a fixed value of KF and different values of nF.
Table 2. Properties of Phenolic Compounds29,30 a phenolic compound
M (g mol-1)
S (g L-1)
S (mmol L-1)
pKa
Tb (K)
phenol 4-nitrophenol 4-chlorophenol 2-chlorophenol
94.1 139.1 128.6 128.6
84.0 15.6 25.5 22.7
893.1 112.1 198.6 176.2
9.92 7.15 9.18 8.49
455 552 493 448
a M ) molar mass; S ) solubility at 298 K; pK ) ionization a constant at 298 K; Tb ) normal boiling temperature.
Table 3. Degree of Ionization at pH 2, 4, 7, and 11 pH phenolic compound PNP OCP PCP P
2
4
0 0 0 0
0 0 0 0
7
11
% of ionization 42 3 1 0
100 100 99 92
H, N, O, and S was carried out using a Eurovector EA elemental analyzer. The concentrations of surface acid and basic sites were determined by back-titrations after equilibrating the samples for 48 h at 298 K with excess 0.01 M NaOH and 0.01 M HNO3, respectively. The points of zero charge were determined using the mass titration method.7,8 The results of these analyses are given in Table 1. The phenolic compounds used were phenol (99+%), 4-nitrophenol (99+%), 4-chlorophenol (99+%), and 2-chlorophenol (99+%), all from Aldrich, and will be designated by P, PNP, PCP, and OCP, respectively. Some relevant properties of the phenolic compounds are given in Table 2. Using the ionization constants and the corresponding ionization constant expressions, it is a simple matter to calculate the percent ionization of each phenolic compound at the different pH values used. These values are given in Table 3. 2. Methods. For the determination of the adsorption isotherms, a fixed amount (∼0.1 g) of each carbon and 50 mL of an aqueous solution of the phenolic compound of desired concentration were placed into a 100 mL stoppered Erlenmeyer. All of the flasks were shaken for 7 days using a thermostated shaker bath at 298 K. This procedure was repeated for different pH values, namely, pH 2 (adjusted by addition of an appropriate amount of dilute nitric acid), pH 4 or pH 7 (the point of zero charge of NSXox and NSXar, respectively), and pH 11 (adjusted by addition of an appropriate amount of dilute sodium hydroxide). After that time, the solutions were filtered and the amount of the phenolic compound was measured using a Hitachi U-3010 UV/visible (5) Carrott, P. J. M.; Roberts, R. A.; Sing, K. S. W. Carbon 1987, 25, 59. (6) Carrott, P. J. M.; Ribeiro Carrott, M. M. L.; Roberts, R. A. Colloids Surf. 1991, 58, 385. (7) Noh, J.; Schwarz, J. Carbon 1990, 28, 675. (8) Carrott, P. J. M.; Ribeiro Carrott, M. M. L.; Candeias, A. J. E.; Ramalho, J. P. P. J. Chem. Soc., Faraday Trans. 1995, 91, 2179.
spectrophotometer at appropriate wavelength. The wavelengths used were 399, 298, 293, and 287 nm for PNP, PCP, OCP, and P, respectively. The determination of each phenolic compound was made in conditions of pH which assured that only one form of the phenolic compound was present. Prior to the analysis of each series of solutions, calibration curves at the same wavelength and pH were determined.
Results and Discussion 1. Modification of the Freundlich Equation. It is necessary to begin by considering the frequently asserted statement that KF and nF (or 1/nF) are measures of adsorption capacity and heterogeneity, respectively. Figure 1 shows adsorption isotherms simulated using the Freundlich Equation with a constant KF and different values of nF. nF is clearly a measure of the curvature of the isotherms and, in this sense, can be considered a measure of the adsorption heterogeneity. A complimentary interpretation is that nF is related to the mean adsorption energy. A low value of nF corresponds to a curved isotherm and hence to a heterogeneous distribution of adsorption energies and a low mean adsorption energy, whereas a high value of nF corresponds to a rectangular isotherm, a more homogeneous distribution of adsorption energies and a higher mean adsorption energy. It is also apparent that the differences are highly significant for low values of nF, but that for nF > 4, the exact value has increasingly less influence on the form of the isotherm. With regard to the parameter KF, its units are dependent on the value of nF and both its units and its numerical value are dependent on the units of Ceq and nads. Most authors have used (nads/mg g-1 + Ceq/mg L-1)9-21 or (nads/ mmol g-1 + Ceq/mmol L-1)22-25 and we will designate the (9) Dutta, S.; Basu, J. K.; Ghar, R. N. Sep. Purif. Technol. 2001, 21, 227. (10) Haghseresht, F.; Lu, G. Q. Energy Fuels 1998, 12, 1100. (11) Le Cloirec, P.; Brasquet, C.; Subrenat, E. Energy Fuels 1997, 11, 331. (12) Shirgaonkar, I. Z.; Joglekar, H. S.; Mundale, V. D.; Joshi, J. B. J. Chem. Eng. Data 1992, 37, 175. (13) Brasquet, C.; Le Cloirec, P. Langmuir 1999, 15, 5906. (14) Jung, M.-W.; Ahn, K.-H.; Lee, Y.; Kim, K.-P.; Rhee, J.-S.; Park, J. T.; Paeng, K.-J. Microchem. J. 2001, 70, 123. (15) Kilduff, J. E.; King, C. J. Ind. Eng. Chem. Res. 1997, 36, 1603. (16) Arafat, H. A.; Franz, M.; Pinto, N. G. Langmuir 1999, 15, 5997. (17) Teng, H.; Hsieh, C.-T. Ind. Eng. Chem. Res. 1998, 37, 3618. (18) Lin, S. H.; Hsu, F. M. Ind. Eng. Chem. Res. 1995, 34, 2110. (19) Tancredi, N.; Medero, N.; Mo¨ller, F.; Pı´riz, J.; Plada, C.; Cordero, T. J. Colloid Interface Sci. 2004, 279, 357. (20) Roostaei, N.; Tezel, F. H. J. Environ. Management 2004, 70, 157. (21) Monneyron, P.; Faur-Brasquet, C.; Sakoda, A.; Suzuki, M.; Le Cloirec, P. Langmuir 2002, 18, 5163. (22) Juang, R.-S.; Wu, F.-C.; Tseng, R.-L. J. Chem. Eng. Data 1996, 41, 487.
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corresponding values of KF by KFg and KFm, respectively. To remove the dependence on the units of nads, it is useful to also consider a parameter KFgm, equal to KFg/M, which corresponds to (nads/mmolg-1 + Ceq/mg L-1). As KF increases the adsorption at a fixed solute equilibrium concentration increases. However, the exact relationship between KF and the overall adsorption capacity will depend on which units are used for Ceq. KFm is the adsorption at Ceq ) 1 mmol L-1, whereas KFgm (and KFg) is the adsorption at Ceq ) 1 mg L-1. If we consider that the molar mass of a simple aromatic solute typically lies in the range 100200, then KFgm (and KFg) evidently corresponds to a point at a much lower solute equilibrium concentration and hence also to a much lower adsorption than KFm. If adsorption is considered as a competitive process in which the adsorbent and the solvent compete for the solute molecules then, in appropriate cases, we can isolate the influence of the adsorbent from that of the solvent by normalizing the adsorption with respect to the solubility of the solute.26,27 Although this was pointed out over 30 years ago, relatively few authors have followed this recommendation, with the notable exception of some recent work involving the DA equation.28 If we do the same when applying the Freundlich equation, then the constant, which we will designate by KFS, will be the adsorption, in the same units as nads, when the equilibrium concentration is equal to the solubility of the solute
( )
nads ) KFS
Ceq S
Figure 2. Representative calculated normalized adsorption isotherms corresponding to P adsorption with nF ) 2 and to PNP adsorption with nF ) 8.
1/nF
(2)
and its value will be independent of both nF and the units of Ceq. This is a very significant advantage of KFS over KFm or KFg. The main disadvantage of this approach is that KFS now refers to a point on the isotherm significantly beyond the normal range of experimental measurements. One possible alternative would be to adopt 1% of the solubility as a reference point. In this case, eq 1 becomes
(
nads ) KF1
)
100Ceq S
1/nF
(3)
where S is the solubility in the same units as Ceq and the constant, now designated by KF1, will also have the same units as nads. We have suggested the value of 1% of the solubility, as this generally falls well within the normal range of experimental measurements. An alternative would be to adopt 10% of the solubility as a reference point
( )
nads ) KF10
10Ceq S
1/nF
(4)
The reference point will now often fall near to the upper limit of the normal range of experimental measurements and KF10 might therefore be a more appropriate parameter to use as an estimate of the relative adsorption capacity of the adsorbent for different solutes. (23) Edgehill, R. U.; Lu, G. Q. J. Chem. Technol. Biotechnol. 1998, 71, 27. (24) Ania, C. O.; Parra, J. B.; Pis, J. J. Fuel Proc. Technol. 2002, 79, 265. (25) Ania, C. O.; Parra, J. B.; Pis, J. J. Fuel Proc. Technol. 2002, 77-78, 337. (26) Mattson, J. S.; Mark, H. B. Activated Carbon; Marcel Dekker: New York, 1971. (27) Lyklema, J. Fundamentals of Interface and Colloid Science; Academic Press: London, 1995; Vol. II. (28) Stoeckli, F.; Lo´pez-Ramo´n, M. V.; Moreno-Castilla, C. Langmuir 2001, 17, 3301.
Figure 3. Variation of KFgm with nF in published data9-23 for (a) P and (b) PNP.
To illustrate the differences that can be observed, Figure 2 shows the localization of the different KF’s on 2 isotherms calculated assuming high and low values of nF and the properties of PNP and P, respectively. The KF value that shows the greatest difference from KFS is KFgm, that is, the most commonly used KF value. When the solubility is high and the isotherm highly curved (for example, P with nF ) 2) KFgm may be only ∼1% of KFS. It is clear that in these cases, in particular, KFgm (and hence KFg) cannot be considered merely as an estimate of the adsorption capacity of the adsorbent. As we will now show, it also appears to be significantly affected by adsorption affinity. It is interesting to note from Figure 1 that the isotherms have a common intersection point at Ceq ) 1 and the numerical value of KF should therefore be independent of nF. As a test, we used published data9-23 for the adsorption of P and PNP and for all KF values, except KFgm (and KFg), we found no correlation between KF and nF. The results obtained with KFgm are shown in Figure 3 from where it can be seen that, despite a certain amount of dispersion of the data, the overall tendency is in each case for KFgm
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Figure 4. Experimental adsorption isotherms at 298 K on Norit SX+ before (NSXar) and after (NSXox) oxidation. P ) phenol ) O; PCP ) 4-chlorophenol ) 3; OCP ) 2-chlorophenol ) 4; PNP ) 4-nitrophenol ) 0.
to increase with nF. This occurs because the reference point of 1 mg L-1 corresponds to only ∼0.001% of the P solubility or ∼0.006% of the PNP solubility. It follows that KFgm is a measure of the initial, almost zero coverage, slope of the isotherm, which, like nF, is strongly influenced by the solute-solid adsorption energy. Hence, by applying the normal form of the Freundlich equation and using units of mg L-1 for Ceq, as is done in the majority of published work, we obtain two parameters providing somewhat similar information, related to adsorption affinity, but no reliable estimate of the overall adsorption capacity of the solid. 2. Application of the Normalized Freundlich Equation. Our experimental adsorption isotherms are given in Figure 4, where the equilibrium concentrations have been normalized with respect to the solubility in pure water. It is clear from Figure 4 that there are some very large differences between the isotherms, which vary from almost linear to highly rectangular. In some cases, most notably P, PNP, and OCP on NSXar at pH 2, there is good agreement between the reduced isotherms. However, in most cases, this is not observed and it is therefore clear that differences in the solubilities of the phenolic
compounds are not sufficient to explain the differences in behavior observed in the various systems studied. Equation 2, in the logarithmic form, was applied to the isotherm data over the approximate range of Ceq/S from 0.001 to 0.13 in order to calculate nF and KFS, and the corresponding values of KF10, KF1, KFm, and KFgm were then calculated. The solid lines in Figure 4 were drawn using the derived Freundlich parameters, and it can be seen that in most cases it was possible to obtain a good fit of the equation to the experimental data despite the large variations in precise isotherm shape. The derived values of nF and of the KF’s are represented in Figure 5. The values of nF and KFm or KFgm are of similar magnitude to those obtained by other authors.9-23 The estimated uncertainty in the calculated nF and KF values based on the standard deviations of the slope and intercept of the least-squares fits was in all cases less than (0.5. It can be seen from Figure 5 that, globally, the most significant difference between the different KF’s is their magnitude which varies in the predicted order KFS > KF10 > KF1 ∼ KFm > KFg. In addition, if we consider just those systems with nF > 4, the relative values on passing from one type of KF to another do not vary very much, as
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Figure 5. Derived values of nF, KFS, KF10, KF1, KFm, and KFgm. Legend on (a) applies to all chartssfor each phenolic compound the order is NSXar @ pH 2, pH 7, pH 11 followed by NSXox @ pH 2, pH 4. Note that scale on (b) is different to (c)-(f)).
predicted above. We believe that the overall agreement between experiment and theory indicates that the mechanism of adsorption is the same in each case and that, for each phenolic compound, the shape of the isotherm (and hence the corresponding nF and KF values) are determined principally by the nature of the surface. On this basis, we can draw a number of conclusions from the data in Figure 5. Protonation of the surface at pH 2 reduces the mean adsorption energy and capacity (lower nF and KF’s). The variation is quite small, but is in agreement with the view that adsorption is most favorable at the pzc.1-3 Oxidation also reduces the mean adsorption energy slightly (for PNP and PCP). The adsorption capacity increases or decreases, depending on the precise system and KF considered. However, even in those cases where it increases, the variation is small. Hence, the results are also in general agreement with other work that has indicated that adsorption of phenolic compounds is more favorable on reduced surfaces.1-3 Deprotonation at pH 11 has a more significant effect, particularly on the adsorption of PNP. A facile explanation of the reduced adsorption at pH 11 is that it is a result of electrostatic repulsion between the negatively charged carbon surface and phenolate ions. However, the decrease in adsorption capacity is much less than the degree of ionization of the phenolic compounds. Furthermore, with PCP nF hardly changes, whereas with PNP and OCP nF
is reduced but remains >4, that is, the interaction with the surface is still relatively strong. There may be two reasons for this apparent inconsistency. In aqueous solution phenolic compounds dissociate according to the equilibrium
ΦOH(aq) S ΦO-(aq) + H+(aq)
(5)
ΦO- will suffer repulsion by the surface and may not adsorb. On the other hand, ΦOH will adsorb and this will shift eq 5 to the left. Hence the effective degree of ionization, considering both ΦOH(aq) and ΦOH(ads) will be less than that indicated in Table 3. We would expect that suppression of the ionization would be most effective for the phenolic compound with the highest pKa, that is, PCP. The results in Figures 4 and 5 confirm that this is the case. On the other hand, we would also expect that suppression of the ionization would be relatively ineffective with PNP as it is 100% ionized at pH 11. It is to be noted, however, that there is still a significant residual adsorption of PNP. This appears to be consistent with the idea that, under any circumstances, there is always a very strong, but kinetically limited, and irreversible (although this was not checked here) adsorption of at least a part of the phenolic compound.1-3 We now turn our attention to those systems which have nF < 4. For these systems, the relative magnitude of KF depends on which one is considered. For instance, if we
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compare P and PNP at pH 7 on NSXar, we find that KFS, KF10, and KF1 are higher for P, whereas KFm and KFgm are lower. If we look at the isotherms in Figure 4c, it is apparent that the use of KFm or KFgm alone could lead to misleading conclusions. We believe that this difference indicates that the mechanism of adsorption is different and that the shape of the isotherms (and hence the corresponding nF and KF values) is influenced not so much by the nature of the surface and its influence on the affinity of the solute for the solid phase, but more by the nature of the solution and its influence on the affinity of the solute for the aqueous phase. This leads us to suggest what we feel is a significant new idea resulting from the work reported here, that is, the notion of a critical adsorption energy, Eca. If the adsorption energy is greater than Eca, then surface control is observed and the form of the reduced adsorption isotherm is determined by the solute-solid interaction energy. On the other hand, if the adsorption energy is less than Eca, then the solution begins to take control of the adsorption process which is then significantly altered. The value of Eca will be dependent on the particular solute in question, and we would not expect (and did not observe) that the adsorption was ever reduced to zero, due to the strong slow irreversible adsorption referred to above. On the other hand, we might expect that, of the four phenolic compounds studied in this work, the one most likely to be influenced by the effect would be phenol. The solubility of phenol is more than three times higher than that of the other compounds and the critical adsorption energy for phenol will therefore be higher. On the other hand, with the exception of OCP, phenol has the lowest boiling point. If we consider that the boiling point is a measure of the strength of the intermolecular forces formed by the molecule, then it might be expected that phenol would have one of the lowest adsorption energies of the compounds studied. In fact, the results show that phenol exhibits characteristics somewhat different to the other compounds. The adsorption isotherms of phenol on NSXar are significantly different to those of all of the other isotherms at pH 7 and pH 11, but are similar at pH 2. Other authors have reported a small step on phenol isotherms which was attributed to a change from horizontal to vertical orientations of the adsorbed molecule.26 Although we did not observe the step, it does seem to be a plausible explanation for the difference in behavior at pH 7 and pH 11 and at pH 2. It is possible that at pH 7 and pH 11 the phenol molecules adsorb vertically (or across the width of the micropores), which would result in a lower solutesolid adsorption energy (and hence lower nF, KFm, and KFgm) but a higher adsorption capacity (and hence higher KFS and KF10). On this basis, KF10 would be the most suitable parameter for characterizing the corresponding adsorption isotherms, namely, in relation to the effective adsorption capacity. At pH 2, the behavior of phenol on NSXar is entirely comparable to that of PNP or PCP on both NSXar and NSXox at all pH values, which would indicate that in this case the phenol molecules adsorb horizontally. Oxidation of the surface has a much more dramatic effect on the adsorption of phenol, and to a certain extent, OCP, than on the adsorption of PNP and PCP. Apparently, the adsorption energy is greater than Eca on NSXar, but the lower adsorption energy on NSXox is less than Eca. In the case of OCP, the adsorption is not reduced to quite the same extent which suggests that the adsorption energy is only slightly less than Eca for OCP. The predominant factor with this molecule appears to be the relatively weak intermolecular forces evidenced by the low boiling point.
Carrott et al.
It is to be noted, however, that despite the weaker solutesolid interaction energy, the overall adsorption capacity does not appear to be altered to a great extent. Once again, in this case KF10 appears to be the best parameter for characterizing the adsorption isotherms. 3. Final Note on Solubility. It might seem reasonable to suggest that the amounts adsorbed ought to be normalized with respect to the solubility at the working pH which in general will be different than the solubility in pure water. On one hand, increasing ionic strength results in a decrease in solubility.30 On the other hand, the solubilities of all of the phenolic compounds studied here are lower in acid solution and higher in basic solution.30 Unfortunately, detailed tables of solubilities of phenolic compounds under a wide range of ionic strengths and pH values do not appear to be available. Furthermore, as already pointed out above, preferential adsorption of the nonionized form suppresses the ionization, and this may also lead to an alteration in the solubility, which would be difficult to take into account quantitatively. At the present time, therefore, there appears to be no choice other than to use the solubilities in pure water. Fortunately, these have been determined by a number of independent authors and are readily available in the literature.29-33 Conclusions The need to normalize aqueous phase adsorption isotherms with respect to the solubility of the solute was pointed out more than 30 years ago.26 However, it is a pratice which has not been followed by the majority of researchers. Our analysis confirms the usefulness of this approach for normalizing the adsorption isotherms and also indicates that misleading conclusions may be reached by applying the usual unnormalized form of the Freundlich equation, in particular, when the units used for the solute equilibrium concentration are mg L-1. We recommend the use of the parameter KF10 as a comparative estimate of the effective adsorption capacity for different solutes. The behavior at lower solute equilibrium concentrations can then be inferred (or calculated, if necessary) from the value of the parameter nF. Analysis of our experimental results also indicates a need to distinguish between two regimes of adsorption, characterized by an adsorption energy, Eads, greater than or less than a critical value, Eca. When Eads > Eca, the shape of the adsorption isotherm is determined by solute-solid interactions. On the other hand, when Eads < Eca solute-solution interactions become increasingly important. Confirmation of this idea would obviously require detailed calorimetric measurements to be carried out. However, it is clear that the best approach for obtaining sound conclusions related to the influence of carbon surface chemistry and solute molecular structure on the adsorption process involves normalizing the data and working in the first regime (Eads > Eca) where the influence of extraneous factors has been eliminated. Finally, essentially for two reasons, we also recommend that values of KFg (but not KFm) continue to be quoted. On one hand, this will facilitate comparison with previous published work. On the other hand, KFg gives a measure (29) Achard, C.; Jaoui, M.; Schwing, M.; Rogalski, M. J. Chem. Eng. Data 1996, 41, 504. (30) Jaoui, M.; Achard, C.; Rogalski, M. J. Chem. Eng. Data 2002, 47, 297. (31) Seidell, A. Solubility of Organic Compounds; Van Nostrand Co.: New York, 1941; Vol. II. (32) Ma, K. C.; Shiu, W. Y.; Mackay, D. J. Chem. Eng. Data 1993, 38, 364. (33) Buchholz, K. D.; Pawliszyn, J. Anal. Chem. 1994, 66, 160.
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of the performance of an activated carbon for removing trace contaminants from very dilute solutions and, in this respect, is still a useful parameter. Acknowledgment. This work was supported by the Fundac¸ a˜o para a Cieˆncia e a Tecnologia (Project No.
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POCTI/CTM/38255/2001) with national and European community funds. The authors are also grateful to the NORIT Company for the provision of the activated carbon sample. LA0520886
