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Comparative study of the adsorption of organic ligands on aluminum oxide by titration calorimetry
Applied Geochemistry', Vo[. 8, pp. 127-139, 1993
0883-2927/93 $6.00 + .00 © 1993 Pergamon Press Ltd
Printed in Great Britain
Comparative study of the adsorption of organic ligands on aluminum oxide by titration calorimetry PIERRE BENOIT,* JANET G. HERINGt a n d WERNER STUMM Institute for Water Resources and Water Pollution Control (EAWAG), Swiss Federal Institute of Technology (ETH), Ziirich, CH-8600, Dfibendorf, Switzerland
(Received 17 January 1992; accepted in revised form 9 July 1992) Abstract--Ligand adsorption on d~-Al203 at pH 8 was examined for a series of organic ligands (aromatic acids, monochlorophenols and aliphatic acids) including both monodentate and bidentate ligands. Adsorption isotherms for the aromatic acids exhibited saturation at high dissolved ligand concentrations; saturation was not observed (over the concentration range examined) for the chlorophenols. Small, though measurable, amounts of heat were evolved on reaction of the aromatic acids, the monochlorophenols and propionate (but not of the longer chain fatty acids) with the oxide surface; overall ligand adsorption reactions were exothermic (AHobs < 0). For adsorption of (partially or fully) protonated ligands, the favorable AHob~was due largely to the exothermic proton transfer reaction between phenolic hydroxyl groups of the ligands and hydroxide ions displaced from the oxide surface. The enthalpy corresponding to the ligand-exchange reaction of surface hydroxyl groups for the various ligands (as fully deprotonated species), AHcorr, appeared to be related to the ligand structure. The surface ligandexchange reaction was more exothermic for the dicarboxylic acid phthalate than for the monocarboxylic acids benzoate or propionate or for salicylate and was endothermic for the chlorophenols.
dissolution of oxide and silicate minerals can be dramatically increased in the presence of bidentate REACTIONS at solid-solution interfaces are of funda- organic ligands (FURRER and STUMM, 1986; STUMM mental importance in the geochemical cycling of and WIELAND, 1990). In contrast, monodentate ortrace elements and in the fate and transport of or- ganic ligands, such as benzoic acid, have little effect ganic compounds in the environment. The properties on mineral dissolution rates and can competitively of mineral surfaces and their reactivity, particularly inhibit dissolution by bidentate ligands (FURRERand toward dissolution, can be markedly influenced by S T U M M , 1986). With natural organic substances (humic substances and plant root exudates), both the sorption of organic compounds. The surface charge of alumium and iron oxides, which are posi- inhibition and acceleration of mineral weathering (depending on pH) have been observed (OcHs et al., tively charged at the pH of natural waters, can be reversed as a result of the adsorption of low molecu1993). The reactivity of organic compounds associated lar weight organic acids, such as phthalic acid (KUMwith mineral surfaces can also differ significantly MERTand STUMM, 1980), fatty acids (YAP etal., 1981), anionic surfactants (WAKAMATSUand FUERSTENAU, from their reactivity in aqueous solution. Organic1968), or fulvic and humic acids (DAVIS and GLOOR, surface associations can either increase degradation 1981; TIPPING, 1981; DAVIS, 1982; LIANG and MOR- rates or protect the adsorbed organic compound GAN, 1990). The negative surface charge of particles against degradation. The former effect is illustrated in natural waters has been attributed to the effect of by the dependence of the anaerobic degradation rates of aromatic azo compounds and halogenated organic coatings (HUNTER and LISS, 1979). Through such an effect on surface charge, the adsorption of ethanes on sediment concentration (WEBER and organic anions can control particle coagulation. Col- WOLFE, 1987; JAFVERT and WOLFE, 1987) and the latter effect is observed in the alkaline hydrolysis of loidal stability can be either increased or decreased some pesticides and reduction of nitrobenzenes by adsorption of fulvic, humic or fatty acids depending on the extent of surface coverage and the result- (MACALADVand WOLFE, 1984; SANDERSand WOLFE, ing surface charge (TIPPING, 1986; JEKEL, 1986; LIANG 1985). When an organic compound promotes the and MORGAN, 1990; AMAL et al., 1992). The rates of reductive dissolution of minerals, organic degradation is an integral part of the dissolution process as illustrated by the oxidation of substituted phenols *Present address: Institut National Agronomique Paris- during the dissolution of manganese oxides (STONE, Grignon, Laboratoire des Sols, F-78850 Thiverval- 1987; ULRICH and STONE, 1989). Thus the effects of Grignon, France. organic-surface associations on reactivity, either of tAuthor to whom correspondence should be addressed: the surface or the organic compound, cannot be Department of Civil Engineering, 4173 Engineering I, University of California, Los Angeles, CA 90024-1593, evaluated without some insight into the nature of the association. U.S.A. INTRODUCTION
127
P. Benoit et al.
128
Table 1. Studies of adsorption enthalpy Sorbate H÷ H÷ H÷ H÷
H~ Cd~-+ Selenite Fluoride, iodate, phosphate, salicylate Cationic copolymers Surfactants Anionic Surfactants Surfactants
Solid Rutile Hematite Rutile Hematite Goethite Rutile AI, Fe and Sn oxides quartz, calcite apatite, fluorite, kaolinite, feldspar AI oxide Rutile Hematite Goethite Goethite Clay Silica Silica Activated carbon Alumina Silica, alumina, bentonite
Conditions 0.02--0.2 M K N O 3 pH 3-9 0.02--0.2 M KNO 3 pH 3-9 0.01-0.1 M NaNO3 pH 4-10 0.001M NaCI pH 2-10
Method
Ref
Calorimetry
(a)
T dependence (surface charge) Calorimetry
(b)
Calorimetry
(d)
Calorimetry
(e,f)
T dependence
(g)
T dependence
(h)
Calorimetry, T dependence Calorimetry
(i)
Calorimetry
(k)
pH 8-9
Calorimetry
(1)
(Not reported)
Calorimetry
(m)
0.001-0. I M NaCI pH 4-10 0.02-0.2 M KNO3 pH 5-8 0.1 M KCI pH 6.7 0.05 M NaNO3 pH 4 0.001 M NaCI pH 7 (Not reported)
(c)
(j)
References: (a) DE KEIZERet al. (1990), (b) FOKKINKet al. (1989), (C) MACHESKYand ANDERSON(1986), (d) WIERERand DOmAS (1988), (e) MACHESKYand Jgcoas (1991a), (f) MACHESKYand JACOBS(1991b), (g) FOKKINKet al. (1990), (h) BALISTR1ERIand CVlou (1987), (i) MACHESKYet al. (1989), (j) DENOYELet al. (1989), (k) PARTYKAet al. (1986), (I) PARTYKAet al. 0989), (m) NOLL(1987).
Several types of interactions may contribute to the association of organic compounds with mineral surfaces. These include chemical interactions with the surface, which may be either specific (surface complex formation) or non-specific (H-bonding or van der Waals attraction), electrostatic interactions with the surface, interactions with the solvent (hydrophobic exclusion), and second-order interactions such as hydrophobic interactions of sorbed organic molecules or co-adsorption (WESTALL, 1987; TIPPING, 1990). When the organic-surface interaction results in modification of properties or reactivity of the surface or the associated organic compounds, these effects can be indicative of the processes contributing to the association. Changes in the spectroscopic properties of adsorbed organic compounds, examined with Fourier transform infra-red (FTIR) spectroscopy (ZELTNER et a l . , 1986; TEJEDORTEJEDOR et a l . , 1990; KUNG and McBRIDE, 1991), electron spin resonance (ESR) spectroscopy (McBRIDE, 1980, 1982), and fluorescence spectroscopy (HERING and STUMM, 1991), have been attributed to formation of inner-sphere surface complexes. More commonly, however, the contributions of various types of interactions have been evaluated on the basis of the observed partitioning of an organic sorbate between aqueous solution and a mineral surface. Diagnostic features of sorption isotherms include the dependence of the extent of sorption on the concentration of the sorbing species and mineral
surface area, on the concentrations of competing or co-adsorbing species, including protons (i.e. pH dependence), and on ionic strength (STUMM et a l . , 1980; SCHINDLERand STUMM,1987; WESTALL,1987; DZOMaAK and MOREL, 1990). Sorption isotherms exhibiting surface saturation and weak ionic strength dependence are characteristic for (anionic) sorbing species that form inner-sphere surface complexes (HAYES et a l . , 1988). Another informative feature of sorption isotherms is their temperature dependence, from which the enthalpy of adsorption can be evaluated. Adsorption enthalpies, which may also be determined by calorimetry, are directly related to the bond strength of the surface species; previous studies of adsorption enthalpies are listed in Table 1. In this paper, the adsorption of a series of organic ligands (aromatic acids, monochlorophenols and aliphatic acids) on aluminum oxide has been studied by titration calorimetry in order to examine the relation between the structure of the sorbing species and adsorption enthalpy.
MATERIALS AND METHODS Reagents
The aluminum oxide used ("Aluminum oxide C", Degussa) consists essentially of ~-AI203 with a primary
Adsorption of organic ligands on aluminium oxide
129
Table 2. Stability constants for acid-base equilibria and surface complex formation* log K
2 (nm)-t
Solution species [HL]/[H+][L-] [H2L]/[H+][HL -] [HL-I/[H+I[L2- l [H2L]/[H+][HL-] [HL-]/[H+][L 2-] [HL]/[H+][L -] [HL]/[H+][L -] [HL]/[H+][L ] [HL I/[H+][L -]
4.2 2.95 5.41 2.97 13.74 4.87 8.53 9.13 9.42
O-AI203
{~-AI~-I~AIOH }[H + ] (-:AIOH}/{-=AIO- }[H + ]
Benzoic acid Phthalic acid
{==-AIL}/{=-AIOH}[HL]
Salicylic acid
{=-AIHL}/{=-AIOH}[H2L]
7.4 10.0 3.7 7.3 2.4 6.0 0.6
Benzoic acid Phthalic acid Salicylic acid Propionic acid 2-chlorophenol 3-chlorophenol 4-chlorophenol
230 229 297 273 273.5 279.5
Surface specws
(-=AIHL }/(-=AIOH} [H2L] {=AIL-}[H +]/{-=AIOH}[H2L] (---ALL- }[H +]/( ~AIOH} [HeL]
* Constants for solution species (for 25°C, 0 ionic strength) from MARTELLand SMITH(1977). Intrinsic constants for surface species (for 22°C, 0.1 M NaC104) from KUMMERTand SXUMM (1980); ¢ Wavelengths used for analysis of organic compounds (see Materials and Methods). particle size of 20 nm and specific surface area of 110 m2/g (FURRER, 1985). The particles were washed with 0.1 M NaOH and rinsed repeatedly with bidistilled water (after the procedure of KUMMERTand STUMM, 1980). The intrinsic acidity constants of the 6-A1203 surface determined previously by KUMMERTand STUMM(1980) are listed in Table 2; the pHzp ~ (zero point of charge) of the oxide is 8.6. Stock suspensions in 0.02 M NaCIO4 at pH 8 were prepared with a particle concentration of 10 g/1. X-ray diffraction spectra of an oxide suspension that had been stored for several years suggested a very slow surface transformation from bayerite (b-AI203) to gibbsite (AI(OH)3); this effect was not observed for suspensions used in the experiments described below (GIOVANOLI, pers. comm). Benzoic, phthalic and salicylic acids (Merck), 2-, 3-, and 4-chlorophenols (Aldrich) and aliphatic acids (propionic (C3), caprylic (C8), and lauric (C12), Fluka) were used without further purification. Acid-base properties of these ligands are also listed in Table 2
Adsorption isotherms The d-AI203 suspensions (10 g/l, pH 8, I = 0.02 M NaCIO4) were prepared in 25 ml polycarbonate centrifuge tubes with ligand concentrations ranging from 0.02 to 2.0 mM. After the tubes were shaken for 18 h at 25°C, the oxide was separated from the supernatant by centrifugation at 14,500 rpm for 45 min. The dissolved ligand concentration in the supernatant was determined by u.v. spectrophotometry for benzoate, phthalate and salicylate. For the chlorophenols (MCP's), the supernatant was acidified to pH 5 and analyzed by HPLC. Adsorbed ligand concentrations were calculated from the difference between the total and dissolved concentrations. The procedure was repeated at 10°C for benzoate and phthalate. Supernatants were also analyzed for dissolved AI by a colorimetric method (Dou6AN and WILSON, 1974). Only slight dissolution of the oxide occurred during the adsorption experiments (dissolved AI < 1/~M). The pH was measured at the beginning and end of each experiment. The pH electrode was calibrated by Gran titration (25°C) thus reported pH values refer to proton concentration rather than activity. Because no buffer was AG 8:2-6
used, slight drifts in pH were observed, but the pH of the suspensions remained between 7.9 and 8.2. The wavelengths used to measure residual concentrations of the aromatic acids and chlorophenols are reported in Table 2. Standard solutions prepared at pH 8 and I = 0.02 M NaClO 4 were used to generate standard curves. Some of the isotherm experiments were repeated and in this case adsorption data are reported as average values.
Modeling of adsorption isotherms Adsorption data for the aromatic acids were modeled using a Langmuir isotherm {---ALL} - KL,,x[L] 1 + K[L]
(1)
where the adsorbed ligand concentration, {---A 1L}, and the maximum adsorption density, Fm~x, are expressed in mol/g and the dissolved ligand concentration, [L] in mol/1 and K is the conditional equilibrium constant for adsorption. The constants, ['max and K, were obtained by non-linear regression. Adsorption data for the chlorophenols were modeled using a Freundlich isotherm; constants were obtained for the linearized equation: log {-~AIL) = log/( + n log [L]
(2)
with units for concentration terms as described above. For the aliphatic acids, adsorption data of ULRICh et al. (1988) were used. In that study, adsorption isotherms were obtained using laC-labeled compounds with the same 6-A120 3 particles but under different conditions (oxide concentration 1 g/l, pH 4, 6, and 10, I = 0.02M NaCIO4). These isotherms were modeled with Freundlich isotherms (as above); portions of the isotherms indicating saturation (at very high ligand concentrations) were not included in the modeling. Constants for adsorption at pH 8 were obtained by averaging the constants for adsorption at pH 6 and pH 10. The assumption of a uniform decrease in ligand adsorption with increasing pH is in accord with the behavior of low molecular weight organic acids (KUMMERT and SrUMM,
P. Benoit et al.
130
Table 3. Summary of THAM calibration data (accepted value: -47.44 kJ/mol) Heat evolved (m J)
A/-/obs (k J/tool)
7520 2350 1175 587,5 293,8 146.9
-48.44 -47.92 -48.02 -48.15 -49.07 -48.32
+ + + + + +
0.62 0.32 0.29 0.58 1.62 1.30
1980). For the longer chain aliphatic acids, this assumption is supported by the observed pH dependence of oleate adsorption onto hematite; the extent of adsorption decreases with increasing pH in the range of the pHzpc of the oxide (YAP et al., 1981). The parameters obtained by fitting the adsorption data with Langmuir or Freundlich isotherms were used to estimate the extent of adsorption for the conditions of the calorimetric experiments (see below), i.e. dissolved and adsorbed ligand concentrations were calculated for a given total ligand concentration. The uncertainty in the adsorbed ligand concentration was estimated by the propagation of the uncertainty in the adsorption constants (K and Fm~x for the data fitted with a Langmuir isotherm and/~ and n for the data fitted with a Freundlich isotherm) where the uncertainty in the dissolved ligand concentration was assumed to be less than or equal to the uncertainty in the adsorbed ligand concentration.
Calorimetric titrations Enthalpy measurements were performed using a TRONAC (Model 450) titration microcalorimeter. All experiments were conducted in an adiabatic mode with an initial suspension volume of 50 ml. The heat evolved in a reaction was measured as the adiabatic temperature change recorded with a thermistor. The conversion to total reaction enthalpy was done after a calibration using a heating resistor (EATHOUGHet al., 1974), In order to test the calorimeter and monitoring software, the heat of neutralization of 50 ml of 0.1 M THAM (Tris(hydroxymethyl)aminomethane) with varying concentrations of HCI was determined. These titrations were carried out so that the total heat production was similar to that in the suspension titrations (as discussed by MACHESKY and JACOBS,1991a). These results are summarized in Table 3. Titrations of suspensions were conducted as follows. Fifty milliliters of 10 g/1 6-A12aO3 suspension was transferred to the calorimetric Dewar reaction vessel after the initial pH had been measured. The pH of the ligand solution (at initial concentrations ranging from 0.001 to 0.1 M) was adjusted to pH 8 by addition of small amounts of solutions of NaOH or HC104. The burette was filled with the ligand solution (titrant). Both titrant and suspension were thermally equilibrated by immersion of the burette and reaction vessel for 30--45 rain in the thermostated (25°C) water bath of the calorimeter. In the titration, 1.6 ml of titrant was delivered continuously by the burette over 3 min and the temperature increase was recorded automatically. Solution blanks were also monitored for heat and pH changes. Except for propionate, heat and pH changes were small (<10 mJ; <0.1 pH units). Non-chemical heat contributions included heat loss through the walls of the Dewar vessel, stirring and heats of dilution. Standard formulas were used to correct for these effects (EATaOUOn et al., 1974). The correction for extraneous chemical reactions was more problematic. Dissolution of the oxide during the calorimetric titrations was assumed to be negligible based on the very limited dissol-
ution observed in the much longer adsorption experiments. Because changes in pH of the suspension during adsorption were also small, adsorption of the added ligand on to the alumina surface was assumed to be the most important chemical process contributing to the measured enthalpies. After correction for non-chemical effects, the experimental data, corresponding to the total heat evolved (Q) during each titration, were converted to enthalpies (AHob~) expressed in kilojoules per mole ligand adsorbed. The number of moles adsorbed in a given titration and the corresponding adsorbed ligand concentrations (F) were estimated from the adsorption isotherms for the various ligands (see above), Then -Q AHob~ - ktmoles adsorbed where Q is the total heat evolved (in m J) and measured enthalpy of reaction (in k J/tool).
(3) AHoo s
iS the
Modeling of surface speciation Distribution of surface species as a function of total ligand concentration was calculated using H Y D R A Q L (PAPELISet al., 1988), a version of MINEQL (WESTALLet al., 1976) that incorporates surface complexation models. The diffuse double layer model was applied to account for electrostatic effects. Intrinsic constants for surface acidity and surface complex formation (see Table 2) and the maximum exchange capacity of surface hydroxyl groups, {~-A1OH }-r, as determined by alkimetric titration (0.21 mol/kg) were taken from KtJM~IERT and STUMM (1980). At pH 8, in the absence of organic ligands, >90% of the surface sites would exist as the neutral hydroxylated species ~-AIOH.
RESULTS AND DISCUSSION
A d s o r p t i o n isotherms In the a d s o r p t i o n i s o t h e r m s for the a r o m a t i c acids p h t h a l a t e , salicylate a n d b e n z o a t e , a p l a t e a u in the a b s o r b e d c o n c e n t r a t i o n was o b s e r v e d at high dissolved c o n c e n t r a t i o n s of t h e ligands (Fig. 1). This b e h a v i o r is consistent with s a t u r a t i o n of the surface by t h e ligand. F o r p h t h a l a t e a n d salicylate, the d a t a agree well with t h e L a n g m u i r i s o t h e r m s (see E q n 1) i n d i c a t e d by the solid (or d a s h e d ) lines in Fig. 1. T h e data for b e n z o a t e m a y also be a p p r o x i m a t e l y described by a L a n g m u i r i s o t h e r m , a l t h o u g h the agreem e n t with the m o d e l is p o o r at low ligand concent r a t i o n s ( c o n s t a n t s are listed in T a b l e 4). A l t h o u g h t h e m a x i m u m a d s o r p t i o n densities o b s e r v e d at p H 8 are quite low, t h e results r e p o r t e d h e r e are consistent with t h e p r e v i o u s studies of KUMMERT a n d STUMM (1980) in which t h e e x t e n t of a r o m a t i c acid adsorption was s h o w n to d e c r e a s e with increasing p H (see below). F o r t h e m o n o c h l o r o p h e n o l s , n o s a t u r a t i o n was o b s e r v e d o v e r the c o n c e n t r a t i o n r a n g e e x a m i n e d (Fig. 2) a n d t h e d a t a were m o d e l e d with a F r e u n d l i c h i s o t h e r m w h e r e the a d s o r b e d ligand c o n c e n t r a t i o n (in tool/g) is e x p r e s s e d as: {---ALL} = I~[L] n
(4)
Adsorption of organic ligands on aluminium oxide
dlich isotherms for the aliphatic acids propionate, caprylate and laurate. The data of ULRICH et al. (1988) for adsorption of these aliphatic acids at pH values of 6 and 10 were fitted with Freundlich isotherms (over the linear region of the isotherms) and values for pH 8 were obtained by interpolation.
30(o) E
20
6 /
131
1
10 8
Effect o f temperature on the adsorption isotherms 0
i
t
i
i
I
0
i
i
i
i
2.0
1.0 dissolved conc. (raM)
30
(b)
20 d 10
8
f i
0
i
,
I
i
i
1.0
2.0
dissolved conc. (raM)
4O 30
~
/ / t /"
6
/
/
20
[]
For phthalate and benzoate, adsorption isotherms were obtained at two temperatures (10 and 25°C). The extent of adsorption was, for both compounds, greater at the lower temperature. The shift in the adsorption isotherms appears to correspond predominantly to a change in the maximum adsorption density; little variation in the adsorption constant is observed (see Table 4). The observed temperature effect on adsorption may be due in part to variation in the pHzpc of the oxide. A shift in the pHzpc of hematite, from 8.6 at 25°C to 9.0 at 10°C, has been reported by FOKKINKet al. (1989). If a similar shift in the pHzpc of aluminium oxide occurs, greater anion adsorption would be expected at lower temperatures (i.e. further below the pHzpc of the oxide at the same pH). In their study of salicylate adsorption onto goethite at pH 4, MACHESKYet al. (1989) found no effect of temperature (from 10 to 40°C) on the adsorption isotherms. These authors suggested, however, that the error inherent in determination of the adsorption isotherms may have obscured the effect of temperature on adsorption.
.9 'd Calorimetric measurements :
0
:
,
:
.
.
.
1.0
.
2.0
dissolved conc. (raM)
FIG. 1. Adsorption isotherms: adsorbed ligand concentrations (pmol/g) as a function of dissolved ligand concentrations (mM) for (a) phthalate, (b) salicylate and (c) benzoate. Conditions: 6-A1203 10 g/l, pH 8, 25°C or 10°C (V in (a),(c)). Solid and dashed lines show model fit to data (Langmuir isotherm; constants in Table 4): ( ) 25°C, ( - - - ) 10°C. (the fitting parameters for the m o d e l , / ( and n, are listed in Table 5). A similar pattern of adsorption of mono-, di- and trichlorophenols on non-crystalline iron oxide (although over a smaller concentration range) has been reported by KUNG and McBmDE (1991). In our study, the extent of adsorption of the monochlorophenols (at pH 8) increases with increasing p K a of the phenols, that is, F4-MCP > F3-MCP > F2.MCP at comparable dissolved ligand concentrations (Fig. 2). This pattern accords with spectroscopic evidence suggesting that the surface bond strength for mono-, di- and trichlorophenols is correlated with the Lewis basicity of the phenolate anions (KoNG and McBRIDE, 1991). Also listed in Table 5 are the constants for Freun-
The calorimetric data for all experiments are summarized in Table 6. The total heat evolved (Q) on titration of aluminum oxide suspensions with the various ligands was greatest for the aromatic acids and least for the aliphatic acids. For caprylate and laurate, no satisfactory measurements could be made; Q < 5 mJ for all ligand concentrations. From the data of ULRICH etal. (1988), we estimated that the adsorption of the aliphatic acids under these experimental conditions should be comparable to the adsorption of the other ligands. Thus, the lack of measurable heat in these titrations is likely to correspond to a very small heat of adsorption rather than to the absence of any adsorption reaction. Adsorption of fatty acids on alumina (through bonding at the carboxylic acid end) has also been demonstrated by ESR spectroscopy of spin probe analogs of stearic (C18) acid (McBRIDE, 1980). For propionic acid, the heat evolved was also quite low (<10 mJ) at low total ligand concentrations but increased significantly at higher total ligand concentrations (Table 6). For the monochlorophenols, the total heat evolved in the titrations ranged from <10 to approximately 100 mJ (Table 6, Fig. 3). The highest values of Q were
132
P. Benoit et al. Table 4. Summary of constants obtained from sorption data modeled with Langmuir isotherms Ligand Salicylate Phthalate Benzoate*
T(°C)
K + 1 S.D.
Fm,x (mol/g) + 1 S.D.
25 25 10 25 10
(1.95 + 0.62) x 10 3 (4.45 + 0.47) x 103 (3.24 + 0.88) x 103 3.7 x 103 1.6 × 103
(3.28 -+ 0.49) x 10 -5 (1.23 _+ 0.04) x 10 -5 (2.07 _+ 0.18) x 10-5 2.1 x 10 5 5.0 x 10 -5
*Error could not be estimated for benzoate. observed for the adsorption of 4-chlorophenol. A t comparable total ligand concentrations, the lowest values of Q were observed for adsorption of 2chlorophenol.
20,
(o)
A 0
.& 8
o
15. o 10 5 0
0.2
0
0.4
0.6
0.8
1.0
0.8
1.0
d[a,=,olved cone. (raM)
(b) 151 0 - ,
20-6
8 £ o
5-
0
0.2
0.4
0.6
dissolved cone. (rnM)
20-
Overall, the highest values of Q were measured for the aromatic acids, particularly salicylic acid (Table 6, Fig. 3). For benzoate, relatively small variation in Q (from 67 to 93 m J) was observed over the experimental range in total ligand concentrations. For phthalate, Q ranged from 66 to 181 mJ and, for salicylate, from 156 to 305 mJ. In Fig. 3, total heat evolved is shown as a function of the adsorbed ligand concentrations as estimated from the adsorption isotherms for the various ligands. For phthalate, salicylate and 4-chlorophenol, there appears to be a trend of increasing heat evolved with increasing moles adsorbed. A quantitative assessment of this trend, however, is not possible because of the uncertainty in the adsorbed ligand concentration. This uncertainty also complicates further interpretation of the calorimetric data because the enthalpies of reaction (AHobs) are calculated from the heat evolved per mole ligand adsorbed. The fairly large uncertainties in the AHohs values shown in Fig. 4 derive from the uncertainties in the adsorption isotherms rather than from the calorimetric measurements. N o n e the less, Fig. 4 indicates that, for all the ligands, the overall reaction with the aluminum oxide suspension is exothermic (AH < 0). Comparison of the values of AHobS for the various ligands, however, requires some consideration of the specific reactions contributing to the measured enthalpies.
C o n t r i b u t i o n s to m e a s u r e d enthalpies o f reaction
(e)
Surface complex formation at the aluminum o x i d e - w a t e r interface involves a surface ligandexchange reaction in which the incoming liquid re-
15d ~o 10-
Table 5. Summary of constants obtained from sorption data modeled with Freundlich isotherms
[]
Ligand 1D O
0
0.2
0.4
0.6 0.8 dissolved cone. (raM)
.0
Fie. 2. Adsorption isotherms: adsorbed ligand concentrations ~mol/g) as a function of dissolved ligand concentrations (mM) for (a) 4-chlorophenol, (b) 3-chlorophenol and (c) 2-chlorophenol. Conditions: 6-A1203 10 g/l, pH 8, 25°C. Solid lines show model fit to data (Freundlich isotherm; constants in Table 5).
4-chlorophenol 3-chlorophenol 2-chlorophenol Propionate* Caprylate Laurate
log I~ + 1 S.D.
n + 1 S.D.
- 1.70 + 0.17 -2.30 + 0.11 -3.09 + 0.21 - 1.7 - 1.3 0.76
1.05 + 0.10 0.89 + 0.06 0.71 + 0,13 0.88 0.97 1.0
*Values for propionate, caprylate, and laurate were obtained by extrapolation from data of ULRICHet al. (1988) and error could not be estimated.
Adsorption of organic ligands on aluminium oxide
133
Table 6. S u m m a r y of calorimetric data
Ligand
pmoles adsorbed + 1 S.D.
[L]T (mM)
Phthalate
0,109 0.143 0.285 1.24 1.55 3.10
1.6 2.0 3.2 5.3 5.5 5.9
Salicylate
0.310 0.465 1/.775 1.55 2.33 3.10
5.0 6.7 9.1 12.2 13.6 14.3
Benzoate *
0.310 //.465 (/.62/) /1.775 1.21 1.55 3.10
4.7 6.0 6.9 7.5 8.5 8.9 9.8
4-MCP
0.031 0.31 0.62 1.05 1.55
(1.2 +_ 0.1 2.0 ± 0.8 4.0 ± 1.5 6.9 -+ 2.5 10.4 ± 3.6
3-MCP
0.31 1/.775 1t.961 1.24 1.55
1.8 4.1 5.0 6.2 7.6
± ± ± ± ±
2-MCP
11.31 0.775 1.24 1.55 2.112
1.3 2.6 3.6 4.2 5.1
± ± ± ± ±
Propionatc ~'
0.155 0.310 1.55 4.03
3 6 26 63
6 12 51 120
O. 16-1.6 0.16-1.6
3-30 8--80
6~0 15-150
Caprylate* Laurate:'
_ + ± + + +
F (umol/g) +1 S.D.
AHob~ (kJ/mol) + 1 S.D.
Q (mJ)
0.2 0.3 0.4 0.8 0.8 0.9
3.2 3.9 6.1 10.3 10.7 11.5
+ _ + + + +
0.4 0.5 0.8 1.5 1.5 1.7
66.2 67,9 105,7 145,0 113,6 180,6
-41 -34 -33 -27 -21 -31
+ + + + + +
± 2.1 ± 2.8 _+ 3.8 +_ 5.2 ± 6.0 ± 6.3
9.6 13.0 17.7 23.6 26.3 27.8
+ 3.8 _+ 5.3 + 7.3 ± 10.1 _+ 11.6 +_ 12.4
197,5 157,1 210,4 265,2 304.8 265.6
-40 -24 -23 -22 -23 -19
+ 16 + 10 _+ 10 +_ 10 _+ 9 ± 8
9.2 12 13 15 16 17 19
67.7 66.9 82.2 66.8 81.1 88.0 92.8
- 14 - 11 - 12 -9 -10 -10 -9
5 4 4 4 3 5
0.4 3.8 7.8 13.5 20.2
+ 0.2 ± 1.5 + 2.9 _+ 4.8 ± 6.9
3.9 36.9 70.7 74.5 103.2
- 2 2 -+ I1 - 1 9 _+ 8 -18 ± 7 -I1 + 4 -10+ 3
0.4 0.9 1.0 1.3 1.5
3.5 7.9 9.6 12.1 14.8
± 0.8 ± 1.7 -+ 2.0 ± 2.4 ± 2.9
29.2 55.6 49.5 32.8 30.0
-16-+ 4 -14 ± 3 - I 0 -+ 2 -5 ± 1 -4 ± 1
0.7 1.2 1.5 1.8 2.1
2.6 5.0 7.0 8.2 9.9
± ± ± ± ±
7.5 10.6 21.6 38.4 11.9
1.3 2.2 3.0 3.4 4.4
-6 -4 -6 -9 -2
± ± ± ± ±
4.0 6.3 38.0 119.7
- 1 - I - 1 -2
<5 <5
---
3 2 3 4 1
:'Error could not be estimated for benzoate, propionate, caprylate or laurate. p l a c e s a s u r f a c e h y d r o x y l g r o u p . A t p H 8, n e a r t h e pHzp ,. o f d - A l 2 0 3 , t h i s r e a c t i o n m a y b e w r i t t e n as: OH
/ ~A1
+ RCO~
\
O~CR -
/ ~
-=AI
+ OH-
\
OH2
OH 2
for reaction with a deprotonated, monodentate l i g a n d ( s u c h as b e n z o a t e o r p r o p i o n a t e ) . R e a c t i o n w i t h a fully d e p r o t o n a t e d , b i d e n t a t e l i g a n d ( s u c h as p h t h a l a t e ) c a n b e w r i t t e n in a s i m i l a r m a n n e r : / -=AI
OH
/ + R ( C O 2 ) 2 ~ ~AI
\ OH 2
In the adsorption of a protonated, monodentate ligand or of a partially protonated, bidentate ligand, h o w e v e r , a d d i t i o n a l r e a c t i o n s c o n t r i b u t e to t h e measured enthalpies. For the reactions of the chlorophenols,
O2C U -
\
/
/ ~AI
OH + CICoH4OH ~ ~AI
\
OC ~.,H4171
/
+ H20
\
OH ~
OH 2
and of salicylate
/ OH
/02C\"~
R + O H - + H20.
O2C
--=AI
+ C~,Ha(CO;-)(OH) ~ ~-AI
\ OH 2
In both of these cases, the measured enthalpy then c o r r e s p o n d s to t h e e n t h a l p y o f t h e l i g a n d - e x c h a n g e r e a c t i o n at t h e s u r f a c e .
\
/
C~,H4 + 2H20
O
the contributions of the enthalpies of the dissociation of the phenolic hydroxyl group and of the neutral-
P. Ben•it et al.
134
contributions to the free energy of reaction are likely to be significant (see below).
(a) 300
A•
200
Enthalpies of surface reactions other than ligand adsorption
(73 100-
0
I 5
0
110
lt5
210
25
/zmoles adsorbed 150(b)
%-,
~
100-
•
.
1/2 AI203 + 3/2 H20 ~ AI(OH)~
~
AH = - 0 . 8 kJ/mol 1/2 A1203 + O H - + 3/2 H20 ~- AI(OH)4
--~-k--
o 50-
0
.
For completeness, the possible contributions of two other reactions, surface protonation and oxide dissolution, to the measured enthalpies should also be considered. The enthalpies of oxide dissolution reactions, in the near neutral to basic pH range, are given by
~lF-II-,~ 0
AH = - 15.5 kJ/mol.
•
t 5
1to
15
,u,males adsorbed
Fla. 3. Heat evolved, Q in mJ, on reaction of ligands with 6A1203 as a function of Mmoles of ligand adsorbed (+1 S.D.) for (a) aromatic acids: (O) phthalate, (A) salicylate, (D) benzoate and (b) monochlorophenols: (O) 4chlorophenol, (A) 3-chlorophenol, (11) 2-chlorophenol. Conditions: 6-A1203 10 g/l, pH 8, 25°C. (Note: error could not be estimated for benzoate.)
Dissolution reactions, however, are relatively slow on the time scale of the calorimetric titrations. Thus any temperature increase due to dissolution reactions would be expected to continue even after titrant addition was complete. That such an effect was not observed indicates that the measured enthalpies were not due to dissolution reactions. This contention is
-60 ¸
ization reaction between the displaced protons and hydroxide ions must be considered. The enthalpies for the acid dissociation reactions are given in Table 7. By accounting for the contributions of these reactions to the measured enthalpies, we may compare the enthalpies corresponding to the ligand-exchange reaction between the surface hydroxyl groups and the various ligands as fully deprotonated species (as shown in Fig. 5); note that the assumption of complete protonation of the phenols at pH 8 may lead to some over-correction in the case of 2-chlorophenol because this phenol is partially deprotonated at the experimental pH. Comparison of Figs 4 and 5 indicates that, for the chlorophenols and salicylate, the favorable AHobS is due to an exothermic proton transfer reaction between the phenolic hydroxyl group and O H - . The values of AHcorr in Fig. 5 suggests that the enthalpy of the surface ligandexchange reaction is related to the structure of the sorbing ligand; the surface ligand-exchange reaction with phthalate is more exothermic than the reaction with benzoate or salicylate and the reactions of the chlorophenols are endothermic. In the overall adsorption reactions of (partially or fully) protonated ligands, the enthalpy associated with proton transfer contributes favorably to the total enthalpy change. Even the overall reactions, however, are not strongly exothermic (Fig. 4) which suggests that entropic
"6
E -40
2
-20 ¸
00
1tO
2tO adsorbed conc.
-40-
3~0
4~0
50
(~mol/g)
(b)
-30
E -20"
-10-
0
0
I
10 adsorbed conc.
210
30
(/,zmol/g)
FIG, 4. Measured enthalpies for reaction of various ligands with 6-A1203 (AHob~in kJ/mol adsorbed +_ 1 S.D.) as a function of adsorbed ligand concentration in Mmol/g (+1 S.D.) for (a) aromatic acids: (O) phthalate, (A) salicylate, (D) benzoate and (b) monochlorophenols: (O) 4chlorophenol, (A) 3-chlorophenol, ( . ) 2-chlorophenol. Conditions: 6-A1203 I0 g/l, pH 8, 25°C. (Note: error could not be estimated for benzoate.)
Adsorption of organic ligands on aluminium oxide Table 7. Heats of acid dissociation* Reaction
AH (kJ/mol)
H20 = H + + OHC6H4(CO2-)(OH) = H + + C6H4(CO2-)OCIC6HaOH = H + + CIC6H404-chlorophenol 3-chlorophenol 2-chlorophenol
56 36 25 25 20
*Reference: CHRISTENSENet al. (1976).
supported by the low concentrations of dissolved A1 ( < 1/~M) measured in the adsorption experiments. The possible contributions of proton adsorption or desorption at the oxide surface to the measured enthapies are more difficult to estimate. Ligandexchange reactions (with fully deprotonated ligands) at the surface displace surface hydroxyl groups and thus increase solution pH, but this pH increase may be buffered by the surface reaction =-A1OH + O H - ~ A I O - + H20 AH -- - 16 kJ/mol (AH determined by MACHESKY and JACOBS, 1991a). To the extent that such a buffering reaction with the surface occurs, the measured enthalpies for reaction of benzoate and phthalate with aluminum oxide would be more negative than the enthalpy of the surface ligand-exchange reaction alone (which would be correspondingly less exothermic than indicated in Fig. 5). The level of surface protonation may also increase in response to the sorption of organic ligands, particularly bidentate ligands for which the charge of the surface species changes with the replacement of the surface hydroxyl group by the ligand. This effect may
135
be illustrated by model calculations for the adsorption of phthalate and salicylate on aluminum oxide as shown in Fig. 6; these calculations employ the intrinsic adsorption constants reported by KUMraERX and SruMM (1980) (see Table 2); note that adsorbed ligand concentrations predicted by the model calculations exceed the observed values. However, model parameters were not optimized for this data set. Proton adsorption is exothermic; values reported by MACHESKYand JACOBS (1991a) range from ----20 kJ/ mol at p H 4 to - 4 5 kJ/mol at pH 9 at an ionic strength of 0.01 M. Thus increase in surface protonation concomitant with sorption of phthalate or salicylate would again result in the measured enthalpy being more negative than the enthalpy of the surface ligand-exchange reaction alone.
Calculated effects of surface coverage on the free energy of adsorption The model calculations for adsorption of bidentate ligands on aluminum oxide also indicate that the apparent stability constant for the surface complex decreases with increasing surface coverage due to the changes in surface charge (or potential). For the reaction ~ A I O H + L 2- + H + ~
---ALL- + H 2 0
--2=AIOH
"5 v
--4 o 't: 6
-6
(o) -8
t
t
,
t
-60
I -5
,
t
i
/
, /i -4
,
log [total phtholote
f
/
/ t =AIHL
/ t
,
,
I -3
q
,
,
J
-2
(M)]
-2-AIOH
-~
-20
.
0
~ ~ r
: v'-
""
,
I
-
'
A
A
~
-4.
T
,5
o -6' 4C
,[ 10
. . . . . . . . . . . . . . . . . 20 30 40 adsorbed
¢0n¢,
/
50
(,u, m a l / c j )
Fro. 5. Enthalpies corresponding to surface ligandexchange reaction of surface hydroxyl groups with various ligands as fully deprotonated species (AHcorr in kJ/mol adsorbed + 1 S.D.) as a function of adsorbed ligand concentrations in/~mol/g (+ 1 S.D.) for (@) phthalate, (A) salicylate, (C]) benzoate, (V) propionate, (V) 4-chlorophenol, (¢) 3-chlorophenol, and (11) 2-chlorophenol. Conditions as Fig. 4. (Note: error could not be estimated for benzoate or propionate.)
(b) -8
l
~ I ,
t
,
I
"~"
,
,
~
,
,
J
,
-4
-5
log [total
solicylate'
t
-.3
J
'
'
'
-2
(M)]
FIG. 6. Surface speciation as calculated with HYDRAQL (with constants in Table 2 and {---AIOH}T -- 0.21 mol/kg): log [concentration of surface species (M)] as a function of log [total ligand concentration (M)] for (a) phthalate and (b) salicylate. Symbols show data from adsorption isotherms.
P. Benoit et al.
136
change in entropy; the reaction is slightly endothermic. The enthalpy of the reaction (AH = 2.7 kJ/ mol) may be calculated from reported values as shown in Table 8. The enthalpy of the corresponding surface ligand-exchange reaction
1.0
0.5
0.0
-=A1OH + L 2- ~-- ---ALL- + O H -0.5
--I
,0
I
-6
'
I
I
I
'
I
'
I
I
,
,
'
-5 -4 log [totol ligand (M)]
'
~
I
-3
I
'
I
-2
FIG. 7. Change in apparent equilibrium constant for ligand adsorption (A log Kapp) as a function of log [total ligand concentration (M)] for phthalate ( - - - ) and salicylate ( ) as calculated with HYDRAQL. Decrease in the apparent constant with increasing total (and thus surface) ligand concentrations is due to electrostatic effects as the surface charge changes due to ligand adsorption.
the relation between the intrinsic and apparent constants is given by:
for salicylate is relatively small, though exothermic (Fig. 5), suggesting that entropic effects may also be important in the formation of surface complexes. The predominance of entropic factors in adsorption of salicylate on goethite at pH 4 has previously been proposed by MACHESKY et al. (1989). The surface complex formation reaction of salicylate with Fe(III) on the surface of goethite is significantly more exothermic (AH ~ - 2 4 kJ/mol at low surface coverage) than the corresponding reaction with Al(III) on the surface of 6-A1203 consistent with the difference in the enthalpy changes of the corresponding reactions in solution (for Fe 3+; A H = - 2 8 . 3 kJ/mol) (MACHESKYet al., 1989).
Kint
Kapp = exp ( - F ~ / R T )
(5) SUMMARY AND CONCLUSIONS
where W is the surface or diffuse layer potential, F is the Faraday constant, and R is the gas constant (F/RT = 38.92/V at 25°C). The predicted variation in the apparent stability constants for adsorption of phthalate and salicylate as a function of the total ligand concentration is shown in Fig. 7. Formation of the surface complex becomes less favorable with increasing total (and surface) ligand concentrations; the corresponding change in AG ° for surface complex formation (over the concentration range of the calorimetric experiments) is +3.7 k J/tool for phthalate and +5.0 kJ/mol for salicylate. Because the effect of surface coverage on adsorption enthalpy is unlikely to be greater than its effect on the free energy, no strong dependence of AH on surface coverage (over this concentration range) would be expected.
Compar&on of complex formation in solution and at the mineral-water interface Complex formation in solution is a ligandexchange reaction in which water molecules (or hydroxide ions) are displaced from the inner coordination sphere of the metal. The release of coordinated waters from the relatively constrained metal complex results in a significant increase in entropy; this effect is particularly important for compact, highly charged cations such as A13÷. The formation of the 1:1 aluminum-salicylate complex by the ligand-exchange reaction A1OH 2+ + L 2- ~- AlL + + O H (L = salicylate) proceeds because of a favorable
The adsorption of organic ligands on the surface of aluminum oxide (at pH 8) involves a surface ligandexchange reaction in which a surface hydroxyl group is displaced by the sorbing ligand. Both mono- and bidentate surface complexes may be formed (depending on the structure of the ligand). Adsorption of (partially or fully) protonated ligands results in a neutralization reaction between protons displaced from the ligand and hydroxide ions displaced from the surface (ScHINDLERand STUMM,1987; WESTALL, 1987). The series of organic ligands studied here includes both monodentate ligands (benzoate, the chlorophenols and the aliphatic acids) and bidentate ligands (salicylate and phthalate) some of which are fully deprotonated at the experimental pH (benzoate, phthalate and the aliphatic acids) and others which are partially or fully protonated (salicylate and the chlorophenols). Adsorption isotherms for the aromatic acids exhibit saturation at high dissolved ligand concentrations whereas no saturation was observed (over the concentrations range examined) for the chlorophenols. Small, though measurable, amounts of heat were evolved on the reaction of propionate, the aromatic acids and the chlorophenols with the &A1203 surface; no satisfactory measurements could be made for the fatty acids caprylate and laurate. Overall, reactions of the aromatic acids, particularly salicylate, resulted in the most heat evolution. The calculation of reaction enthalpies from the measured heat evolved was subject to considerable uncertainty because of the necessity of estimating adsorbed ligand concentrations from adsorption isotherms. None the less, examination of the reaction enthalpies may
Adsorption of organic ligands on aluminium oxide
137
Table 8. Enthalpies and free energies of complexation reactions in solution* Reaction
AH (kJ/mol)
AI3+ + HL = AlL + + H + H + + L2- = HLA1OH2-" = AI3+ + OHA1OH2+ + L2- = AlL ÷ + OH-
32.6 -36 +6.1 +2.7
log Kt
AG° (kJ/mol)
0.183 13.74 -9.01 4.91
-1.05 -78.43 51.43 -28.05
* L = salicylate; references: CHR1STENSENet al. (1976); CHRISTENSENand IzA-rr (1983); MACHESKYand JACOaS(1991a); MARTELLand SMITH(1977); t Stability constants for reactions as written.
provide some insight into processes occurring at the mineral surface. The overall reactions of the sorbing ligands with the 6-A1203 surface were exothermic for all ligands (i.e. other than the fatty acids). The relatively small values of AHobs, however, suggest that favorable entropy changes may be important (perhaps even predominant) in the formation of surface complexes. For reaction of protonated ligands (salicylate and the chlorophenols), the favorable enthalpies of reaction (i.e. AHobs) may be due largely to the exothermic proton transfer reaction between the phenolic hydroxy group and hydroxide ions displaced from the oxide surface. If the enthalpic contributions of such reactions are accounted for, the enthalpies of the surface ligand-exchange reaction (AHcorr), that is for the reaction -=AIOH + L"- ~- =-ALL~'-1)- + O H (n = 1 or 2) may be compared for the various ligands. The values of AHcorr appear to be related to the chelating functionality of the ligand; the surface ligand-exchange reaction is more exothermic for phthalate than for benzoate or salicylate and endothermic for the chlorophenols. It should be noted, however, that the surface ligand-exchange reactions would be less exothermic than the values of AHcorr would indicate if other (exothermic) reactions, particularly surface protonation, had contributed to the measured enthalpies. The calculation of AHco~r allows some estimation of the contribution of the surface ligand--exchange reaction to the overall adsorption enthalpy. It is the enthalpy of the surface ligand-exchange reaction that should be most directly related to the structure and bond strength of the surface complex. A pronounced influence of the structure of surface complexes on surface reactivity toward dissolution has been demonstrated for aluminum oxide (FURRER and STUMM, 1986). The extent of A1 and P release from soil mediated by organic acids has been shown to vary with the specific organic acid used (Fox et al., 1990). Thus the mobilization of A1 and of substances, such as P, that adsorb strongly to aluminum oxides from soils will be influenced not only by the concentration but also by the composition of dissolved organic matter in soil waters.
The relation of surface structure and reactivity has implications for the fate and transport of surfaceassociated organics as well as for mineral weathering. The transport of organic substances, including pollutants such as pesticides, in soils and aquifers will be retarded by their adsorption onto immobile solid surfaces. Hydrolysis, an important mechanism for the degradation of thiophosphate pesticides, has been shown to be catalyzed by oxide surfaces when the structure of the organic ester undergoing hydrolysis allows for formation of a bidentate surface c o m p l e x (TORRENTS and STONE, 1991). Thus both the extent and the nature of organic-surface interactions will affect the persistence of organic compounds in the environment. Acknowledgements--We thank Professor R. Giovanoli
(University of Bern) for sample analysis and Professor I. Grenth6 (Royal Institute of Technology. Stockholm) for helpful discussions at early stages of this work. The participation of P. Benoit was sponsored by the Institut National de la Recherche Agronomique, Thiverval-Grignon, France. This work was supported in part by the Swiss National Science Foundation. Editorial handling: J. G. Catts.
REFERENCES
AMAL R., RAPERJ. A. and WAITET. D. (1992) Effect of fulvic acid adsorption on the aggregation kinetics and structure of hematite particles. J. Colloid Interface Sci. 151,244-257. BALASTRIERIL. S. and CHou T. T. (1987) Selenium adsorption by goethite. Soil Sci. Soc. Am. J. 51, 1145-1151. CHRISTENSENJ. J., HANSENL. D. and IZAYrR. M. (1976) Handbook of Proton Ionization Heats. Wileylnterscience. CHRISTENSEN J. J. and IZATr R. M. (1983) Handbook of Metal Ligand Heats. Marcel Dekker. DAVISJ. A. (1982) Adsorption of natural dissolved organic matter at the oxide/water interface. Geochim. Cosmochim. Acta 46, 2381-2393. DAVIS J. A. and GLOOR R. (1981) Adsorption of dissolved
organics in lake water by aluminum oxide. Effect of molecular weight. Envir. Sci. Tech. 15, 1223-1229. DENOYEL R., DURAND G., LAFUMA F. and AUDEBERT R.
(1989) Adsorption of cationic polyelectrolytes onto montmorillonite and silica: a microcalorimetric study of their conformation. J. Colloid Interface Sci. 139,281-290. DOUGANW. K. and WILSONg. L. (1974) Absorptiometric determination of aluminum in water. Comparison of
138
P. Benoit et al.
some chromogenic reagents and the development of an improved method. Analyst 99,413-430. DZOMBAK D. A. and MOREL F. M. M. (1990) Surface Complexation Modeling: Hydrous Ferric Oxide. WileyInterscience. EATHOUGH D. J., CHRISTENSENJ. J. and IZATr R. M. (1974) Experiments in Thermometric Titrimetry and Titration Calorimetry. Brigham Young University Press. FOKKINK L. G. J., DE KEIZER A. and LYKLEMAJ. (1989) Temperature dependence of the electrical double layer on oxides: Rutile and hematite. J. Colloid Interface Sci. 127, 116-131. FOKKINK L. G. J., DE KEIZER A. and LYKLEMAJ. (1990) Temperature dependence of cadmium adsorption on oxides. J. Colloid Interface Sci. 135, 118-131. Fox T. R., COMERFORD N. B. and MCFEE W. W. (1990) Phosphorous and aluminum release from a spodic horizon mediated by organic acids. Soil Sci. Soc. Am. J. 54, 1763-1767. FURRER G. (1985) Die oberflaechenkontrollierte Aufloesung von Metalloxiden: ein koordinationschemischer Ansatz zur Verwitterungskinetik. Ph.D. Thesis, Swiss Federal Institute of Technology, Zurich. FURRER G. and STUMMW. (1986) The coordination chemistry of weathering: I. Dissolution kinetics of 6-AbeO3 and BeO. Geochim. cosmochim. Acta 50, 1847-1860. HAYES K. F., PAPELISC. and LECKIEJ. O. (1988) Modeling ionic strength effects on anion adsorption at hydrous oxide/solution interfaces. J. Colloid Interface Sci. 125, 717-726. HER1NG J. G. and STurau W. (1991) Fluorescence spectroscopic evidence for surface complex formation at the mineral-water interface: elucidation of the mechanism of ligand-promoted dissolution. Langmuir 7, 1567-1570. HUNTER K. A. and LIss P. S. (1979) The surface charge of suspended particles in estuarine and coastal waters. Nature 282,823-825. JAFVERT C. T. and WOLFE N. U (1987) Degradation of selected halogenated ethanes in anoxic sediment-water systems. Envir. Toxicol. Chem. 6, 827-837. JEKEL M. R. (1986) The stabilization of dispersed mineral particles by adsorption of humic substances. Water Res. 12, 1543-1554. DE KEIZER A., FOKKINK L. G. J. and LYKLEMAJ. (1990) Thermodynamics of proton charge formation on oxides. Microcalorimetry. Colloids Surf. 49,149-163. KUMMERTR. and STUMraW. (1980) The surface complexation of organic acids on hydrous 7-A1203. J. Colloid Interface Sci. 75,373-385. KUNG K.-H.S. and MCBRIDE M. B. (1991) Bonding of chlorophenols on iron and aluminum oxides. Envir. Sci. Tech. 25,702-709. LIANG L. and MORGANJ. J. (1990) Chemical aspects of iron oxide coagulation in water: laboratory studies and implications for natural systems. Aquatic Sci. 52, 32-55. MACALADV D. L. and WOLFE N. L. (1984) Abiotic hydrolysis of sorbed pesticides. In Treatment and Disposal of Pesticide Wastes (eds R. F. KRUEGER and J. N. SEIaER), pp. 221--244, ACS Syrup. Ser. 259. MACHESKYM. L. and ANDERSON M. A. (1986) Calorimetric acid-base titrations of aqueous goethite and futile suspensions. Langmuir 2, 582-587. MACHESKY M. L., BISCHOFF B. L. and ANDERSON M. A. (1989) Calorimetric investigation of anion adsorption onto goethite. Envir. Sci. Tech. 23, 580-587. MACnESRV M. L. and JACOBSP. F. (1991a) Titration calorimetry of aqueous alumina suspensions Part I. Results and comparisons with similar studies. Colloids Surf. 53,297314. MACHESKYM. L. and JACOBSP. F. (1991b) Titration calorimetry of aqueous alumina suspensions Part II. Discussion of enthalpy changes with pH and ionic strength. Colloids Surf. 53,315-328.
MARTELL A. E. and SMITh R. M. (1977) Critical Stability Constants, Vol. 3. Plenum Press. MCBRtDE M. B. (1980) Adsorption of fatty acid spin probes on amorphous alumina. J. Colloid Interface Sci. 76,393398. MCBRIDE M. B. (1982) Organic anion adsorption on aluminum hydroxides: Spin probe studies. Clay Clay Minerals 30,438-444. NOEL L. A. (1987) Adsorption calorimetry of surfactant interaction with minerals. Colloids Surf. 26, 43-54. OCHS M., BRUNNER I., STUMMW. and t~osowC B. (1993) Effects of root exudates and humic substances on weathering kinetics. Water Air Soil Pollut. (in press). PAPEEIS C., HAYES K. F. and LECKIE J. O. (1988) HYDRAQL: a program for the computation of chemical equilibrium composition of aqueous batch systems including surface complexation modeling of ion adsorption at the oxide/solution interface. Tech. Rep. No. 306. Dept. of Civil Engineering. Stanford Univ., Stanford, California. PARTYKAS., LINDHEIMERM., ZA1NIS., KEH E. and BRUN B. (1986) Improved calorimetric method to investigate adsorption processes from solution onto solid surfaces. Langmuir 2, 101-105. PARTYKA S., RUDZUNKA W., BRUN B. and CLINT J. H. (1989) Calorimetric studies of adsorption of anionic surfactants onto alumina. Langmuir, 5,297-304. SANDERS P. and WOLF N. L. (1985) Reduction of nitroaromatic compounds in anaerobic, sterile sediments. Presented at the National Meeting of the Am. Chem. Soc., Chicago, Illinois. SCHINDLERP. W. and STUMMW. (1987) The surface chemistry of oxides, hydroxides, and oxide minerals. In Aquatic Surface Chemistry (ed. W. STUMM), pp. 83--110. WileyInterscience. STONE A. T. (1987) Reductive dissolution of manganese (Ill/IV) oxides by substituted phenols. Envir. Sci. Tech. 21,979-988. STUMM W., KUMMERT R. and SIGG L. (1980) A ligand exchange model for the adsorption of inorganic and organic ligands at hydrous oxide interfaces, Croatica Chem. Acta 53,291-312. STUMMW. and WIELANDE. (1990) Dissolution of oxide and silicate minerals: Rates depend on surface speciation. In Aquatic Chemical Kinetics (ed. W. STUMM),pp. 367-400. Wiley-lnterscience. TEJEDOR-TEJEDOR M. I., YOST E. C. and ANDERSONM. A. (1990) Characterization of benzoic and phenolic complexes at the goethite/aqueous solution interface using cylindrical internal reflection Fourier transform infrared spectroscopy. Part I. Methodology. Langmuir 6, 979987. TIPPING E. (1981) The adsorption of aquatic humic substances by iron oxides. Geochim. Cosmochim. Acta 45, 191-199. TIPPING E. (1986) Some aspects of the interactions between particulate oxides and aquatic humic substances. Mar. Chem. 18,161-169. TIPPING E. (1990) Interactions of organic acids with inorganic and organic surfaces. In Organic Acids in Aquatic Ecosystems (ed. E. M. PERDUEand E. T. (GJESSING),pp. 209--221. John Wiley and Sons. TORRENTS A. and STONE A. T. (1991) Hydrolysis of phenyl picolinate at the mineral/water interface. Envir. Sci. Tech. 25,143-149. ULRICH H.-J. and STONE A. T. (1989) Oxidation of chlorophenols adsorbed to manganese oxide surfaces. Envir. Sci. Tech. 23,421-428. ULRICH H.-J., STUMMW. and (~osovI~ B. (1988) Adsorption of aliphatic fatty acids on aquatic interfaces. Comparison between two model surfaces: the mercury electrode and ~-A120 3 colloids. Envir. Sci. Tech. 22, 37-41. WAKAMATSUT. and FEURSTENAUD. W. (1968) The effect of
Adsorption of organic ligands on aluminium oxide hydrocarbon chain length on the adsorption of surfactants at the solid/water interface. In Adsorption from Aqueous Solution (ed. W. I. WEBER and E. MATEJWlC), pp. 161--172. ACS Adv. Chem. Ser. 79. WEBER E. J. and WOLFE N. L. (1978) Kinetic studies of the reduction of aromatic azo compounds in anaerobic sediment/water systems. Envir. Toxicol. Chem. 6, 911919. WESTALL J. C. (1987) Adsorption mechanisms in aquatic surface chemistry. In Aquatic Surface Chemistry (ed. W. STUMM)pp. 3--32. Wiley-lnterscience. WESXALLJ. C., ZACHARYJ. L. and MORELF. M. M. (1976) MINEQL: a computer program for the calculation of the chemical equilibrium composition of aqueous systems, R. M. Parsons Laboratory Tech. Note No. 18, Massachusetts Institute of Technology, Cambridge.
139
WIERERK. A. and DOBIASB. (1988) Exchange enthalpies of H + and O H - adsorption on minerals with different characters of potential-determining ions. J. Colloid Interface Sci. 122, 171-177. YAPS. N., M1SHRAR. K., RAGHAVENS. and FUERSTENAUD. W. (1981) The adsorption of oleic acid on goethite. In Adsorption from Aqueous Solutions (ed. P. H. TEWARI), pp. 119--142. Plenum Press. ZELTNER W. A., YOST E. C., MACHESKYM. L., TEJEDORTEJEDOR M. I. and ANDERSONM. A. (1986) Characterization of anion binding on goethite using titration calorimetry and cylindrical internal reflection-Fourier transform infrared spectroscopy. In Geochemical Processes at Mineral Surfaces (ed. J. A. DAVIESand K. F. HAYES),pp. 142--161. ACS Symp. Ser. 323.
0883-2927/93 $6.00 + .00 © 1993 Pergamon Press Ltd
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Comparative study of the adsorption of organic ligands on aluminum oxide by titration calorimetry PIERRE BENOIT,* JANET G. HERINGt a n d WERNER STUMM Institute for Water Resources and Water Pollution Control (EAWAG), Swiss Federal Institute of Technology (ETH), Ziirich, CH-8600, Dfibendorf, Switzerland
(Received 17 January 1992; accepted in revised form 9 July 1992) Abstract--Ligand adsorption on d~-Al203 at pH 8 was examined for a series of organic ligands (aromatic acids, monochlorophenols and aliphatic acids) including both monodentate and bidentate ligands. Adsorption isotherms for the aromatic acids exhibited saturation at high dissolved ligand concentrations; saturation was not observed (over the concentration range examined) for the chlorophenols. Small, though measurable, amounts of heat were evolved on reaction of the aromatic acids, the monochlorophenols and propionate (but not of the longer chain fatty acids) with the oxide surface; overall ligand adsorption reactions were exothermic (AHobs < 0). For adsorption of (partially or fully) protonated ligands, the favorable AHob~was due largely to the exothermic proton transfer reaction between phenolic hydroxyl groups of the ligands and hydroxide ions displaced from the oxide surface. The enthalpy corresponding to the ligand-exchange reaction of surface hydroxyl groups for the various ligands (as fully deprotonated species), AHcorr, appeared to be related to the ligand structure. The surface ligandexchange reaction was more exothermic for the dicarboxylic acid phthalate than for the monocarboxylic acids benzoate or propionate or for salicylate and was endothermic for the chlorophenols.
dissolution of oxide and silicate minerals can be dramatically increased in the presence of bidentate REACTIONS at solid-solution interfaces are of funda- organic ligands (FURRER and STUMM, 1986; STUMM mental importance in the geochemical cycling of and WIELAND, 1990). In contrast, monodentate ortrace elements and in the fate and transport of or- ganic ligands, such as benzoic acid, have little effect ganic compounds in the environment. The properties on mineral dissolution rates and can competitively of mineral surfaces and their reactivity, particularly inhibit dissolution by bidentate ligands (FURRERand toward dissolution, can be markedly influenced by S T U M M , 1986). With natural organic substances (humic substances and plant root exudates), both the sorption of organic compounds. The surface charge of alumium and iron oxides, which are posi- inhibition and acceleration of mineral weathering (depending on pH) have been observed (OcHs et al., tively charged at the pH of natural waters, can be reversed as a result of the adsorption of low molecu1993). The reactivity of organic compounds associated lar weight organic acids, such as phthalic acid (KUMwith mineral surfaces can also differ significantly MERTand STUMM, 1980), fatty acids (YAP etal., 1981), anionic surfactants (WAKAMATSUand FUERSTENAU, from their reactivity in aqueous solution. Organic1968), or fulvic and humic acids (DAVIS and GLOOR, surface associations can either increase degradation 1981; TIPPING, 1981; DAVIS, 1982; LIANG and MOR- rates or protect the adsorbed organic compound GAN, 1990). The negative surface charge of particles against degradation. The former effect is illustrated in natural waters has been attributed to the effect of by the dependence of the anaerobic degradation rates of aromatic azo compounds and halogenated organic coatings (HUNTER and LISS, 1979). Through such an effect on surface charge, the adsorption of ethanes on sediment concentration (WEBER and organic anions can control particle coagulation. Col- WOLFE, 1987; JAFVERT and WOLFE, 1987) and the latter effect is observed in the alkaline hydrolysis of loidal stability can be either increased or decreased some pesticides and reduction of nitrobenzenes by adsorption of fulvic, humic or fatty acids depending on the extent of surface coverage and the result- (MACALADVand WOLFE, 1984; SANDERSand WOLFE, ing surface charge (TIPPING, 1986; JEKEL, 1986; LIANG 1985). When an organic compound promotes the and MORGAN, 1990; AMAL et al., 1992). The rates of reductive dissolution of minerals, organic degradation is an integral part of the dissolution process as illustrated by the oxidation of substituted phenols *Present address: Institut National Agronomique Paris- during the dissolution of manganese oxides (STONE, Grignon, Laboratoire des Sols, F-78850 Thiverval- 1987; ULRICH and STONE, 1989). Thus the effects of Grignon, France. organic-surface associations on reactivity, either of tAuthor to whom correspondence should be addressed: the surface or the organic compound, cannot be Department of Civil Engineering, 4173 Engineering I, University of California, Los Angeles, CA 90024-1593, evaluated without some insight into the nature of the association. U.S.A. INTRODUCTION
127
P. Benoit et al.
128
Table 1. Studies of adsorption enthalpy Sorbate H÷ H÷ H÷ H÷
H~ Cd~-+ Selenite Fluoride, iodate, phosphate, salicylate Cationic copolymers Surfactants Anionic Surfactants Surfactants
Solid Rutile Hematite Rutile Hematite Goethite Rutile AI, Fe and Sn oxides quartz, calcite apatite, fluorite, kaolinite, feldspar AI oxide Rutile Hematite Goethite Goethite Clay Silica Silica Activated carbon Alumina Silica, alumina, bentonite
Conditions 0.02--0.2 M K N O 3 pH 3-9 0.02--0.2 M KNO 3 pH 3-9 0.01-0.1 M NaNO3 pH 4-10 0.001M NaCI pH 2-10
Method
Ref
Calorimetry
(a)
T dependence (surface charge) Calorimetry
(b)
Calorimetry
(d)
Calorimetry
(e,f)
T dependence
(g)
T dependence
(h)
Calorimetry, T dependence Calorimetry
(i)
Calorimetry
(k)
pH 8-9
Calorimetry
(1)
(Not reported)
Calorimetry
(m)
0.001-0. I M NaCI pH 4-10 0.02-0.2 M KNO3 pH 5-8 0.1 M KCI pH 6.7 0.05 M NaNO3 pH 4 0.001 M NaCI pH 7 (Not reported)
(c)
(j)
References: (a) DE KEIZERet al. (1990), (b) FOKKINKet al. (1989), (C) MACHESKYand ANDERSON(1986), (d) WIERERand DOmAS (1988), (e) MACHESKYand Jgcoas (1991a), (f) MACHESKYand JACOBS(1991b), (g) FOKKINKet al. (1990), (h) BALISTR1ERIand CVlou (1987), (i) MACHESKYet al. (1989), (j) DENOYELet al. (1989), (k) PARTYKAet al. (1986), (I) PARTYKAet al. 0989), (m) NOLL(1987).
Several types of interactions may contribute to the association of organic compounds with mineral surfaces. These include chemical interactions with the surface, which may be either specific (surface complex formation) or non-specific (H-bonding or van der Waals attraction), electrostatic interactions with the surface, interactions with the solvent (hydrophobic exclusion), and second-order interactions such as hydrophobic interactions of sorbed organic molecules or co-adsorption (WESTALL, 1987; TIPPING, 1990). When the organic-surface interaction results in modification of properties or reactivity of the surface or the associated organic compounds, these effects can be indicative of the processes contributing to the association. Changes in the spectroscopic properties of adsorbed organic compounds, examined with Fourier transform infra-red (FTIR) spectroscopy (ZELTNER et a l . , 1986; TEJEDORTEJEDOR et a l . , 1990; KUNG and McBRIDE, 1991), electron spin resonance (ESR) spectroscopy (McBRIDE, 1980, 1982), and fluorescence spectroscopy (HERING and STUMM, 1991), have been attributed to formation of inner-sphere surface complexes. More commonly, however, the contributions of various types of interactions have been evaluated on the basis of the observed partitioning of an organic sorbate between aqueous solution and a mineral surface. Diagnostic features of sorption isotherms include the dependence of the extent of sorption on the concentration of the sorbing species and mineral
surface area, on the concentrations of competing or co-adsorbing species, including protons (i.e. pH dependence), and on ionic strength (STUMM et a l . , 1980; SCHINDLERand STUMM,1987; WESTALL,1987; DZOMaAK and MOREL, 1990). Sorption isotherms exhibiting surface saturation and weak ionic strength dependence are characteristic for (anionic) sorbing species that form inner-sphere surface complexes (HAYES et a l . , 1988). Another informative feature of sorption isotherms is their temperature dependence, from which the enthalpy of adsorption can be evaluated. Adsorption enthalpies, which may also be determined by calorimetry, are directly related to the bond strength of the surface species; previous studies of adsorption enthalpies are listed in Table 1. In this paper, the adsorption of a series of organic ligands (aromatic acids, monochlorophenols and aliphatic acids) on aluminum oxide has been studied by titration calorimetry in order to examine the relation between the structure of the sorbing species and adsorption enthalpy.
MATERIALS AND METHODS Reagents
The aluminum oxide used ("Aluminum oxide C", Degussa) consists essentially of ~-AI203 with a primary
Adsorption of organic ligands on aluminium oxide
129
Table 2. Stability constants for acid-base equilibria and surface complex formation* log K
2 (nm)-t
Solution species [HL]/[H+][L-] [H2L]/[H+][HL -] [HL-I/[H+I[L2- l [H2L]/[H+][HL-] [HL-]/[H+][L 2-] [HL]/[H+][L -] [HL]/[H+][L -] [HL]/[H+][L ] [HL I/[H+][L -]
4.2 2.95 5.41 2.97 13.74 4.87 8.53 9.13 9.42
O-AI203
{~-AI~-I~AIOH }[H + ] (-:AIOH}/{-=AIO- }[H + ]
Benzoic acid Phthalic acid
{==-AIL}/{=-AIOH}[HL]
Salicylic acid
{=-AIHL}/{=-AIOH}[H2L]
7.4 10.0 3.7 7.3 2.4 6.0 0.6
Benzoic acid Phthalic acid Salicylic acid Propionic acid 2-chlorophenol 3-chlorophenol 4-chlorophenol
230 229 297 273 273.5 279.5
Surface specws
(-=AIHL }/(-=AIOH} [H2L] {=AIL-}[H +]/{-=AIOH}[H2L] (---ALL- }[H +]/( ~AIOH} [HeL]
* Constants for solution species (for 25°C, 0 ionic strength) from MARTELLand SMITH(1977). Intrinsic constants for surface species (for 22°C, 0.1 M NaC104) from KUMMERTand SXUMM (1980); ¢ Wavelengths used for analysis of organic compounds (see Materials and Methods). particle size of 20 nm and specific surface area of 110 m2/g (FURRER, 1985). The particles were washed with 0.1 M NaOH and rinsed repeatedly with bidistilled water (after the procedure of KUMMERTand STUMM, 1980). The intrinsic acidity constants of the 6-A1203 surface determined previously by KUMMERTand STUMM(1980) are listed in Table 2; the pHzp ~ (zero point of charge) of the oxide is 8.6. Stock suspensions in 0.02 M NaCIO4 at pH 8 were prepared with a particle concentration of 10 g/1. X-ray diffraction spectra of an oxide suspension that had been stored for several years suggested a very slow surface transformation from bayerite (b-AI203) to gibbsite (AI(OH)3); this effect was not observed for suspensions used in the experiments described below (GIOVANOLI, pers. comm). Benzoic, phthalic and salicylic acids (Merck), 2-, 3-, and 4-chlorophenols (Aldrich) and aliphatic acids (propionic (C3), caprylic (C8), and lauric (C12), Fluka) were used without further purification. Acid-base properties of these ligands are also listed in Table 2
Adsorption isotherms The d-AI203 suspensions (10 g/l, pH 8, I = 0.02 M NaCIO4) were prepared in 25 ml polycarbonate centrifuge tubes with ligand concentrations ranging from 0.02 to 2.0 mM. After the tubes were shaken for 18 h at 25°C, the oxide was separated from the supernatant by centrifugation at 14,500 rpm for 45 min. The dissolved ligand concentration in the supernatant was determined by u.v. spectrophotometry for benzoate, phthalate and salicylate. For the chlorophenols (MCP's), the supernatant was acidified to pH 5 and analyzed by HPLC. Adsorbed ligand concentrations were calculated from the difference between the total and dissolved concentrations. The procedure was repeated at 10°C for benzoate and phthalate. Supernatants were also analyzed for dissolved AI by a colorimetric method (Dou6AN and WILSON, 1974). Only slight dissolution of the oxide occurred during the adsorption experiments (dissolved AI < 1/~M). The pH was measured at the beginning and end of each experiment. The pH electrode was calibrated by Gran titration (25°C) thus reported pH values refer to proton concentration rather than activity. Because no buffer was AG 8:2-6
used, slight drifts in pH were observed, but the pH of the suspensions remained between 7.9 and 8.2. The wavelengths used to measure residual concentrations of the aromatic acids and chlorophenols are reported in Table 2. Standard solutions prepared at pH 8 and I = 0.02 M NaClO 4 were used to generate standard curves. Some of the isotherm experiments were repeated and in this case adsorption data are reported as average values.
Modeling of adsorption isotherms Adsorption data for the aromatic acids were modeled using a Langmuir isotherm {---ALL} - KL,,x[L] 1 + K[L]
(1)
where the adsorbed ligand concentration, {---A 1L}, and the maximum adsorption density, Fm~x, are expressed in mol/g and the dissolved ligand concentration, [L] in mol/1 and K is the conditional equilibrium constant for adsorption. The constants, ['max and K, were obtained by non-linear regression. Adsorption data for the chlorophenols were modeled using a Freundlich isotherm; constants were obtained for the linearized equation: log {-~AIL) = log/( + n log [L]
(2)
with units for concentration terms as described above. For the aliphatic acids, adsorption data of ULRICh et al. (1988) were used. In that study, adsorption isotherms were obtained using laC-labeled compounds with the same 6-A120 3 particles but under different conditions (oxide concentration 1 g/l, pH 4, 6, and 10, I = 0.02M NaCIO4). These isotherms were modeled with Freundlich isotherms (as above); portions of the isotherms indicating saturation (at very high ligand concentrations) were not included in the modeling. Constants for adsorption at pH 8 were obtained by averaging the constants for adsorption at pH 6 and pH 10. The assumption of a uniform decrease in ligand adsorption with increasing pH is in accord with the behavior of low molecular weight organic acids (KUMMERT and SrUMM,
P. Benoit et al.
130
Table 3. Summary of THAM calibration data (accepted value: -47.44 kJ/mol) Heat evolved (m J)
A/-/obs (k J/tool)
7520 2350 1175 587,5 293,8 146.9
-48.44 -47.92 -48.02 -48.15 -49.07 -48.32
+ + + + + +
0.62 0.32 0.29 0.58 1.62 1.30
1980). For the longer chain aliphatic acids, this assumption is supported by the observed pH dependence of oleate adsorption onto hematite; the extent of adsorption decreases with increasing pH in the range of the pHzpc of the oxide (YAP et al., 1981). The parameters obtained by fitting the adsorption data with Langmuir or Freundlich isotherms were used to estimate the extent of adsorption for the conditions of the calorimetric experiments (see below), i.e. dissolved and adsorbed ligand concentrations were calculated for a given total ligand concentration. The uncertainty in the adsorbed ligand concentration was estimated by the propagation of the uncertainty in the adsorption constants (K and Fm~x for the data fitted with a Langmuir isotherm and/~ and n for the data fitted with a Freundlich isotherm) where the uncertainty in the dissolved ligand concentration was assumed to be less than or equal to the uncertainty in the adsorbed ligand concentration.
Calorimetric titrations Enthalpy measurements were performed using a TRONAC (Model 450) titration microcalorimeter. All experiments were conducted in an adiabatic mode with an initial suspension volume of 50 ml. The heat evolved in a reaction was measured as the adiabatic temperature change recorded with a thermistor. The conversion to total reaction enthalpy was done after a calibration using a heating resistor (EATHOUGHet al., 1974), In order to test the calorimeter and monitoring software, the heat of neutralization of 50 ml of 0.1 M THAM (Tris(hydroxymethyl)aminomethane) with varying concentrations of HCI was determined. These titrations were carried out so that the total heat production was similar to that in the suspension titrations (as discussed by MACHESKY and JACOBS,1991a). These results are summarized in Table 3. Titrations of suspensions were conducted as follows. Fifty milliliters of 10 g/1 6-A12aO3 suspension was transferred to the calorimetric Dewar reaction vessel after the initial pH had been measured. The pH of the ligand solution (at initial concentrations ranging from 0.001 to 0.1 M) was adjusted to pH 8 by addition of small amounts of solutions of NaOH or HC104. The burette was filled with the ligand solution (titrant). Both titrant and suspension were thermally equilibrated by immersion of the burette and reaction vessel for 30--45 rain in the thermostated (25°C) water bath of the calorimeter. In the titration, 1.6 ml of titrant was delivered continuously by the burette over 3 min and the temperature increase was recorded automatically. Solution blanks were also monitored for heat and pH changes. Except for propionate, heat and pH changes were small (<10 mJ; <0.1 pH units). Non-chemical heat contributions included heat loss through the walls of the Dewar vessel, stirring and heats of dilution. Standard formulas were used to correct for these effects (EATaOUOn et al., 1974). The correction for extraneous chemical reactions was more problematic. Dissolution of the oxide during the calorimetric titrations was assumed to be negligible based on the very limited dissol-
ution observed in the much longer adsorption experiments. Because changes in pH of the suspension during adsorption were also small, adsorption of the added ligand on to the alumina surface was assumed to be the most important chemical process contributing to the measured enthalpies. After correction for non-chemical effects, the experimental data, corresponding to the total heat evolved (Q) during each titration, were converted to enthalpies (AHob~) expressed in kilojoules per mole ligand adsorbed. The number of moles adsorbed in a given titration and the corresponding adsorbed ligand concentrations (F) were estimated from the adsorption isotherms for the various ligands (see above), Then -Q AHob~ - ktmoles adsorbed where Q is the total heat evolved (in m J) and measured enthalpy of reaction (in k J/tool).
(3) AHoo s
iS the
Modeling of surface speciation Distribution of surface species as a function of total ligand concentration was calculated using H Y D R A Q L (PAPELISet al., 1988), a version of MINEQL (WESTALLet al., 1976) that incorporates surface complexation models. The diffuse double layer model was applied to account for electrostatic effects. Intrinsic constants for surface acidity and surface complex formation (see Table 2) and the maximum exchange capacity of surface hydroxyl groups, {~-A1OH }-r, as determined by alkimetric titration (0.21 mol/kg) were taken from KtJM~IERT and STUMM (1980). At pH 8, in the absence of organic ligands, >90% of the surface sites would exist as the neutral hydroxylated species ~-AIOH.
RESULTS AND DISCUSSION
A d s o r p t i o n isotherms In the a d s o r p t i o n i s o t h e r m s for the a r o m a t i c acids p h t h a l a t e , salicylate a n d b e n z o a t e , a p l a t e a u in the a b s o r b e d c o n c e n t r a t i o n was o b s e r v e d at high dissolved c o n c e n t r a t i o n s of t h e ligands (Fig. 1). This b e h a v i o r is consistent with s a t u r a t i o n of the surface by t h e ligand. F o r p h t h a l a t e a n d salicylate, the d a t a agree well with t h e L a n g m u i r i s o t h e r m s (see E q n 1) i n d i c a t e d by the solid (or d a s h e d ) lines in Fig. 1. T h e data for b e n z o a t e m a y also be a p p r o x i m a t e l y described by a L a n g m u i r i s o t h e r m , a l t h o u g h the agreem e n t with the m o d e l is p o o r at low ligand concent r a t i o n s ( c o n s t a n t s are listed in T a b l e 4). A l t h o u g h t h e m a x i m u m a d s o r p t i o n densities o b s e r v e d at p H 8 are quite low, t h e results r e p o r t e d h e r e are consistent with t h e p r e v i o u s studies of KUMMERT a n d STUMM (1980) in which t h e e x t e n t of a r o m a t i c acid adsorption was s h o w n to d e c r e a s e with increasing p H (see below). F o r t h e m o n o c h l o r o p h e n o l s , n o s a t u r a t i o n was o b s e r v e d o v e r the c o n c e n t r a t i o n r a n g e e x a m i n e d (Fig. 2) a n d t h e d a t a were m o d e l e d with a F r e u n d l i c h i s o t h e r m w h e r e the a d s o r b e d ligand c o n c e n t r a t i o n (in tool/g) is e x p r e s s e d as: {---ALL} = I~[L] n
(4)
Adsorption of organic ligands on aluminium oxide
dlich isotherms for the aliphatic acids propionate, caprylate and laurate. The data of ULRICH et al. (1988) for adsorption of these aliphatic acids at pH values of 6 and 10 were fitted with Freundlich isotherms (over the linear region of the isotherms) and values for pH 8 were obtained by interpolation.
30(o) E
20
6 /
131
1
10 8
Effect o f temperature on the adsorption isotherms 0
i
t
i
i
I
0
i
i
i
i
2.0
1.0 dissolved conc. (raM)
30
(b)
20 d 10
8
f i
0
i
,
I
i
i
1.0
2.0
dissolved conc. (raM)
4O 30
~
/ / t /"
6
/
/
20
[]
For phthalate and benzoate, adsorption isotherms were obtained at two temperatures (10 and 25°C). The extent of adsorption was, for both compounds, greater at the lower temperature. The shift in the adsorption isotherms appears to correspond predominantly to a change in the maximum adsorption density; little variation in the adsorption constant is observed (see Table 4). The observed temperature effect on adsorption may be due in part to variation in the pHzpc of the oxide. A shift in the pHzpc of hematite, from 8.6 at 25°C to 9.0 at 10°C, has been reported by FOKKINKet al. (1989). If a similar shift in the pHzpc of aluminium oxide occurs, greater anion adsorption would be expected at lower temperatures (i.e. further below the pHzpc of the oxide at the same pH). In their study of salicylate adsorption onto goethite at pH 4, MACHESKYet al. (1989) found no effect of temperature (from 10 to 40°C) on the adsorption isotherms. These authors suggested, however, that the error inherent in determination of the adsorption isotherms may have obscured the effect of temperature on adsorption.
.9 'd Calorimetric measurements :
0
:
,
:
.
.
.
1.0
.
2.0
dissolved conc. (raM)
FIG. 1. Adsorption isotherms: adsorbed ligand concentrations (pmol/g) as a function of dissolved ligand concentrations (mM) for (a) phthalate, (b) salicylate and (c) benzoate. Conditions: 6-A1203 10 g/l, pH 8, 25°C or 10°C (V in (a),(c)). Solid and dashed lines show model fit to data (Langmuir isotherm; constants in Table 4): ( ) 25°C, ( - - - ) 10°C. (the fitting parameters for the m o d e l , / ( and n, are listed in Table 5). A similar pattern of adsorption of mono-, di- and trichlorophenols on non-crystalline iron oxide (although over a smaller concentration range) has been reported by KUNG and McBmDE (1991). In our study, the extent of adsorption of the monochlorophenols (at pH 8) increases with increasing p K a of the phenols, that is, F4-MCP > F3-MCP > F2.MCP at comparable dissolved ligand concentrations (Fig. 2). This pattern accords with spectroscopic evidence suggesting that the surface bond strength for mono-, di- and trichlorophenols is correlated with the Lewis basicity of the phenolate anions (KoNG and McBRIDE, 1991). Also listed in Table 5 are the constants for Freun-
The calorimetric data for all experiments are summarized in Table 6. The total heat evolved (Q) on titration of aluminum oxide suspensions with the various ligands was greatest for the aromatic acids and least for the aliphatic acids. For caprylate and laurate, no satisfactory measurements could be made; Q < 5 mJ for all ligand concentrations. From the data of ULRICH etal. (1988), we estimated that the adsorption of the aliphatic acids under these experimental conditions should be comparable to the adsorption of the other ligands. Thus, the lack of measurable heat in these titrations is likely to correspond to a very small heat of adsorption rather than to the absence of any adsorption reaction. Adsorption of fatty acids on alumina (through bonding at the carboxylic acid end) has also been demonstrated by ESR spectroscopy of spin probe analogs of stearic (C18) acid (McBRIDE, 1980). For propionic acid, the heat evolved was also quite low (<10 mJ) at low total ligand concentrations but increased significantly at higher total ligand concentrations (Table 6). For the monochlorophenols, the total heat evolved in the titrations ranged from <10 to approximately 100 mJ (Table 6, Fig. 3). The highest values of Q were
132
P. Benoit et al. Table 4. Summary of constants obtained from sorption data modeled with Langmuir isotherms Ligand Salicylate Phthalate Benzoate*
T(°C)
K + 1 S.D.
Fm,x (mol/g) + 1 S.D.
25 25 10 25 10
(1.95 + 0.62) x 10 3 (4.45 + 0.47) x 103 (3.24 + 0.88) x 103 3.7 x 103 1.6 × 103
(3.28 -+ 0.49) x 10 -5 (1.23 _+ 0.04) x 10 -5 (2.07 _+ 0.18) x 10-5 2.1 x 10 5 5.0 x 10 -5
*Error could not be estimated for benzoate. observed for the adsorption of 4-chlorophenol. A t comparable total ligand concentrations, the lowest values of Q were observed for adsorption of 2chlorophenol.
20,
(o)
A 0
.& 8
o
15. o 10 5 0
0.2
0
0.4
0.6
0.8
1.0
0.8
1.0
d[a,=,olved cone. (raM)
(b) 151 0 - ,
20-6
8 £ o
5-
0
0.2
0.4
0.6
dissolved cone. (rnM)
20-
Overall, the highest values of Q were measured for the aromatic acids, particularly salicylic acid (Table 6, Fig. 3). For benzoate, relatively small variation in Q (from 67 to 93 m J) was observed over the experimental range in total ligand concentrations. For phthalate, Q ranged from 66 to 181 mJ and, for salicylate, from 156 to 305 mJ. In Fig. 3, total heat evolved is shown as a function of the adsorbed ligand concentrations as estimated from the adsorption isotherms for the various ligands. For phthalate, salicylate and 4-chlorophenol, there appears to be a trend of increasing heat evolved with increasing moles adsorbed. A quantitative assessment of this trend, however, is not possible because of the uncertainty in the adsorbed ligand concentration. This uncertainty also complicates further interpretation of the calorimetric data because the enthalpies of reaction (AHobs) are calculated from the heat evolved per mole ligand adsorbed. The fairly large uncertainties in the AHohs values shown in Fig. 4 derive from the uncertainties in the adsorption isotherms rather than from the calorimetric measurements. N o n e the less, Fig. 4 indicates that, for all the ligands, the overall reaction with the aluminum oxide suspension is exothermic (AH < 0). Comparison of the values of AHobS for the various ligands, however, requires some consideration of the specific reactions contributing to the measured enthalpies.
C o n t r i b u t i o n s to m e a s u r e d enthalpies o f reaction
(e)
Surface complex formation at the aluminum o x i d e - w a t e r interface involves a surface ligandexchange reaction in which the incoming liquid re-
15d ~o 10-
Table 5. Summary of constants obtained from sorption data modeled with Freundlich isotherms
[]
Ligand 1D O
0
0.2
0.4
0.6 0.8 dissolved cone. (raM)
.0
Fie. 2. Adsorption isotherms: adsorbed ligand concentrations ~mol/g) as a function of dissolved ligand concentrations (mM) for (a) 4-chlorophenol, (b) 3-chlorophenol and (c) 2-chlorophenol. Conditions: 6-A1203 10 g/l, pH 8, 25°C. Solid lines show model fit to data (Freundlich isotherm; constants in Table 5).
4-chlorophenol 3-chlorophenol 2-chlorophenol Propionate* Caprylate Laurate
log I~ + 1 S.D.
n + 1 S.D.
- 1.70 + 0.17 -2.30 + 0.11 -3.09 + 0.21 - 1.7 - 1.3 0.76
1.05 + 0.10 0.89 + 0.06 0.71 + 0,13 0.88 0.97 1.0
*Values for propionate, caprylate, and laurate were obtained by extrapolation from data of ULRICHet al. (1988) and error could not be estimated.
Adsorption of organic ligands on aluminium oxide
133
Table 6. S u m m a r y of calorimetric data
Ligand
pmoles adsorbed + 1 S.D.
[L]T (mM)
Phthalate
0,109 0.143 0.285 1.24 1.55 3.10
1.6 2.0 3.2 5.3 5.5 5.9
Salicylate
0.310 0.465 1/.775 1.55 2.33 3.10
5.0 6.7 9.1 12.2 13.6 14.3
Benzoate *
0.310 //.465 (/.62/) /1.775 1.21 1.55 3.10
4.7 6.0 6.9 7.5 8.5 8.9 9.8
4-MCP
0.031 0.31 0.62 1.05 1.55
(1.2 +_ 0.1 2.0 ± 0.8 4.0 ± 1.5 6.9 -+ 2.5 10.4 ± 3.6
3-MCP
0.31 1/.775 1t.961 1.24 1.55
1.8 4.1 5.0 6.2 7.6
± ± ± ± ±
2-MCP
11.31 0.775 1.24 1.55 2.112
1.3 2.6 3.6 4.2 5.1
± ± ± ± ±
Propionatc ~'
0.155 0.310 1.55 4.03
3 6 26 63
6 12 51 120
O. 16-1.6 0.16-1.6
3-30 8--80
6~0 15-150
Caprylate* Laurate:'
_ + ± + + +
F (umol/g) +1 S.D.
AHob~ (kJ/mol) + 1 S.D.
Q (mJ)
0.2 0.3 0.4 0.8 0.8 0.9
3.2 3.9 6.1 10.3 10.7 11.5
+ _ + + + +
0.4 0.5 0.8 1.5 1.5 1.7
66.2 67,9 105,7 145,0 113,6 180,6
-41 -34 -33 -27 -21 -31
+ + + + + +
± 2.1 ± 2.8 _+ 3.8 +_ 5.2 ± 6.0 ± 6.3
9.6 13.0 17.7 23.6 26.3 27.8
+ 3.8 _+ 5.3 + 7.3 ± 10.1 _+ 11.6 +_ 12.4
197,5 157,1 210,4 265,2 304.8 265.6
-40 -24 -23 -22 -23 -19
+ 16 + 10 _+ 10 +_ 10 _+ 9 ± 8
9.2 12 13 15 16 17 19
67.7 66.9 82.2 66.8 81.1 88.0 92.8
- 14 - 11 - 12 -9 -10 -10 -9
5 4 4 4 3 5
0.4 3.8 7.8 13.5 20.2
+ 0.2 ± 1.5 + 2.9 _+ 4.8 ± 6.9
3.9 36.9 70.7 74.5 103.2
- 2 2 -+ I1 - 1 9 _+ 8 -18 ± 7 -I1 + 4 -10+ 3
0.4 0.9 1.0 1.3 1.5
3.5 7.9 9.6 12.1 14.8
± 0.8 ± 1.7 -+ 2.0 ± 2.4 ± 2.9
29.2 55.6 49.5 32.8 30.0
-16-+ 4 -14 ± 3 - I 0 -+ 2 -5 ± 1 -4 ± 1
0.7 1.2 1.5 1.8 2.1
2.6 5.0 7.0 8.2 9.9
± ± ± ± ±
7.5 10.6 21.6 38.4 11.9
1.3 2.2 3.0 3.4 4.4
-6 -4 -6 -9 -2
± ± ± ± ±
4.0 6.3 38.0 119.7
- 1 - I - 1 -2
<5 <5
---
3 2 3 4 1
:'Error could not be estimated for benzoate, propionate, caprylate or laurate. p l a c e s a s u r f a c e h y d r o x y l g r o u p . A t p H 8, n e a r t h e pHzp ,. o f d - A l 2 0 3 , t h i s r e a c t i o n m a y b e w r i t t e n as: OH
/ ~A1
+ RCO~
\
O~CR -
/ ~
-=AI
+ OH-
\
OH2
OH 2
for reaction with a deprotonated, monodentate l i g a n d ( s u c h as b e n z o a t e o r p r o p i o n a t e ) . R e a c t i o n w i t h a fully d e p r o t o n a t e d , b i d e n t a t e l i g a n d ( s u c h as p h t h a l a t e ) c a n b e w r i t t e n in a s i m i l a r m a n n e r : / -=AI
OH
/ + R ( C O 2 ) 2 ~ ~AI
\ OH 2
In the adsorption of a protonated, monodentate ligand or of a partially protonated, bidentate ligand, h o w e v e r , a d d i t i o n a l r e a c t i o n s c o n t r i b u t e to t h e measured enthalpies. For the reactions of the chlorophenols,
O2C U -
\
/
/ ~AI
OH + CICoH4OH ~ ~AI
\
OC ~.,H4171
/
+ H20
\
OH ~
OH 2
and of salicylate
/ OH
/02C\"~
R + O H - + H20.
O2C
--=AI
+ C~,Ha(CO;-)(OH) ~ ~-AI
\ OH 2
In both of these cases, the measured enthalpy then c o r r e s p o n d s to t h e e n t h a l p y o f t h e l i g a n d - e x c h a n g e r e a c t i o n at t h e s u r f a c e .
\
/
C~,H4 + 2H20
O
the contributions of the enthalpies of the dissociation of the phenolic hydroxyl group and of the neutral-
P. Ben•it et al.
134
contributions to the free energy of reaction are likely to be significant (see below).
(a) 300
A•
200
Enthalpies of surface reactions other than ligand adsorption
(73 100-
0
I 5
0
110
lt5
210
25
/zmoles adsorbed 150(b)
%-,
~
100-
•
.
1/2 AI203 + 3/2 H20 ~ AI(OH)~
~
AH = - 0 . 8 kJ/mol 1/2 A1203 + O H - + 3/2 H20 ~- AI(OH)4
--~-k--
o 50-
0
.
For completeness, the possible contributions of two other reactions, surface protonation and oxide dissolution, to the measured enthalpies should also be considered. The enthalpies of oxide dissolution reactions, in the near neutral to basic pH range, are given by
~lF-II-,~ 0
AH = - 15.5 kJ/mol.
•
t 5
1to
15
,u,males adsorbed
Fla. 3. Heat evolved, Q in mJ, on reaction of ligands with 6A1203 as a function of Mmoles of ligand adsorbed (+1 S.D.) for (a) aromatic acids: (O) phthalate, (A) salicylate, (D) benzoate and (b) monochlorophenols: (O) 4chlorophenol, (A) 3-chlorophenol, (11) 2-chlorophenol. Conditions: 6-A1203 10 g/l, pH 8, 25°C. (Note: error could not be estimated for benzoate.)
Dissolution reactions, however, are relatively slow on the time scale of the calorimetric titrations. Thus any temperature increase due to dissolution reactions would be expected to continue even after titrant addition was complete. That such an effect was not observed indicates that the measured enthalpies were not due to dissolution reactions. This contention is
-60 ¸
ization reaction between the displaced protons and hydroxide ions must be considered. The enthalpies for the acid dissociation reactions are given in Table 7. By accounting for the contributions of these reactions to the measured enthalpies, we may compare the enthalpies corresponding to the ligand-exchange reaction between the surface hydroxyl groups and the various ligands as fully deprotonated species (as shown in Fig. 5); note that the assumption of complete protonation of the phenols at pH 8 may lead to some over-correction in the case of 2-chlorophenol because this phenol is partially deprotonated at the experimental pH. Comparison of Figs 4 and 5 indicates that, for the chlorophenols and salicylate, the favorable AHobS is due to an exothermic proton transfer reaction between the phenolic hydroxyl group and O H - . The values of AHcorr in Fig. 5 suggests that the enthalpy of the surface ligandexchange reaction is related to the structure of the sorbing ligand; the surface ligand-exchange reaction with phthalate is more exothermic than the reaction with benzoate or salicylate and the reactions of the chlorophenols are endothermic. In the overall adsorption reactions of (partially or fully) protonated ligands, the enthalpy associated with proton transfer contributes favorably to the total enthalpy change. Even the overall reactions, however, are not strongly exothermic (Fig. 4) which suggests that entropic
"6
E -40
2
-20 ¸
00
1tO
2tO adsorbed conc.
-40-
3~0
4~0
50
(~mol/g)
(b)
-30
E -20"
-10-
0
0
I
10 adsorbed conc.
210
30
(/,zmol/g)
FIG, 4. Measured enthalpies for reaction of various ligands with 6-A1203 (AHob~in kJ/mol adsorbed +_ 1 S.D.) as a function of adsorbed ligand concentration in Mmol/g (+1 S.D.) for (a) aromatic acids: (O) phthalate, (A) salicylate, (D) benzoate and (b) monochlorophenols: (O) 4chlorophenol, (A) 3-chlorophenol, ( . ) 2-chlorophenol. Conditions: 6-A1203 I0 g/l, pH 8, 25°C. (Note: error could not be estimated for benzoate.)
Adsorption of organic ligands on aluminium oxide Table 7. Heats of acid dissociation* Reaction
AH (kJ/mol)
H20 = H + + OHC6H4(CO2-)(OH) = H + + C6H4(CO2-)OCIC6HaOH = H + + CIC6H404-chlorophenol 3-chlorophenol 2-chlorophenol
56 36 25 25 20
*Reference: CHRISTENSENet al. (1976).
supported by the low concentrations of dissolved A1 ( < 1/~M) measured in the adsorption experiments. The possible contributions of proton adsorption or desorption at the oxide surface to the measured enthapies are more difficult to estimate. Ligandexchange reactions (with fully deprotonated ligands) at the surface displace surface hydroxyl groups and thus increase solution pH, but this pH increase may be buffered by the surface reaction =-A1OH + O H - ~ A I O - + H20 AH -- - 16 kJ/mol (AH determined by MACHESKY and JACOBS, 1991a). To the extent that such a buffering reaction with the surface occurs, the measured enthalpies for reaction of benzoate and phthalate with aluminum oxide would be more negative than the enthalpy of the surface ligand-exchange reaction alone (which would be correspondingly less exothermic than indicated in Fig. 5). The level of surface protonation may also increase in response to the sorption of organic ligands, particularly bidentate ligands for which the charge of the surface species changes with the replacement of the surface hydroxyl group by the ligand. This effect may
135
be illustrated by model calculations for the adsorption of phthalate and salicylate on aluminum oxide as shown in Fig. 6; these calculations employ the intrinsic adsorption constants reported by KUMraERX and SruMM (1980) (see Table 2); note that adsorbed ligand concentrations predicted by the model calculations exceed the observed values. However, model parameters were not optimized for this data set. Proton adsorption is exothermic; values reported by MACHESKYand JACOBS (1991a) range from ----20 kJ/ mol at p H 4 to - 4 5 kJ/mol at pH 9 at an ionic strength of 0.01 M. Thus increase in surface protonation concomitant with sorption of phthalate or salicylate would again result in the measured enthalpy being more negative than the enthalpy of the surface ligand-exchange reaction alone.
Calculated effects of surface coverage on the free energy of adsorption The model calculations for adsorption of bidentate ligands on aluminum oxide also indicate that the apparent stability constant for the surface complex decreases with increasing surface coverage due to the changes in surface charge (or potential). For the reaction ~ A I O H + L 2- + H + ~
---ALL- + H 2 0
--2=AIOH
"5 v
--4 o 't: 6
-6
(o) -8
t
t
,
t
-60
I -5
,
t
i
/
, /i -4
,
log [total phtholote
f
/
/ t =AIHL
/ t
,
,
I -3
q
,
,
J
-2
(M)]
-2-AIOH
-~
-20
.
0
~ ~ r
: v'-
""
,
I
-
'
A
A
~
-4.
T
,5
o -6' 4C
,[ 10
. . . . . . . . . . . . . . . . . 20 30 40 adsorbed
¢0n¢,
/
50
(,u, m a l / c j )
Fro. 5. Enthalpies corresponding to surface ligandexchange reaction of surface hydroxyl groups with various ligands as fully deprotonated species (AHcorr in kJ/mol adsorbed + 1 S.D.) as a function of adsorbed ligand concentrations in/~mol/g (+ 1 S.D.) for (@) phthalate, (A) salicylate, (C]) benzoate, (V) propionate, (V) 4-chlorophenol, (¢) 3-chlorophenol, and (11) 2-chlorophenol. Conditions as Fig. 4. (Note: error could not be estimated for benzoate or propionate.)
(b) -8
l
~ I ,
t
,
I
"~"
,
,
~
,
,
J
,
-4
-5
log [total
solicylate'
t
-.3
J
'
'
'
-2
(M)]
FIG. 6. Surface speciation as calculated with HYDRAQL (with constants in Table 2 and {---AIOH}T -- 0.21 mol/kg): log [concentration of surface species (M)] as a function of log [total ligand concentration (M)] for (a) phthalate and (b) salicylate. Symbols show data from adsorption isotherms.
P. Benoit et al.
136
change in entropy; the reaction is slightly endothermic. The enthalpy of the reaction (AH = 2.7 kJ/ mol) may be calculated from reported values as shown in Table 8. The enthalpy of the corresponding surface ligand-exchange reaction
1.0
0.5
0.0
-=A1OH + L 2- ~-- ---ALL- + O H -0.5
--I
,0
I
-6
'
I
I
I
'
I
'
I
I
,
,
'
-5 -4 log [totol ligand (M)]
'
~
I
-3
I
'
I
-2
FIG. 7. Change in apparent equilibrium constant for ligand adsorption (A log Kapp) as a function of log [total ligand concentration (M)] for phthalate ( - - - ) and salicylate ( ) as calculated with HYDRAQL. Decrease in the apparent constant with increasing total (and thus surface) ligand concentrations is due to electrostatic effects as the surface charge changes due to ligand adsorption.
the relation between the intrinsic and apparent constants is given by:
for salicylate is relatively small, though exothermic (Fig. 5), suggesting that entropic effects may also be important in the formation of surface complexes. The predominance of entropic factors in adsorption of salicylate on goethite at pH 4 has previously been proposed by MACHESKY et al. (1989). The surface complex formation reaction of salicylate with Fe(III) on the surface of goethite is significantly more exothermic (AH ~ - 2 4 kJ/mol at low surface coverage) than the corresponding reaction with Al(III) on the surface of 6-A1203 consistent with the difference in the enthalpy changes of the corresponding reactions in solution (for Fe 3+; A H = - 2 8 . 3 kJ/mol) (MACHESKYet al., 1989).
Kint
Kapp = exp ( - F ~ / R T )
(5) SUMMARY AND CONCLUSIONS
where W is the surface or diffuse layer potential, F is the Faraday constant, and R is the gas constant (F/RT = 38.92/V at 25°C). The predicted variation in the apparent stability constants for adsorption of phthalate and salicylate as a function of the total ligand concentration is shown in Fig. 7. Formation of the surface complex becomes less favorable with increasing total (and surface) ligand concentrations; the corresponding change in AG ° for surface complex formation (over the concentration range of the calorimetric experiments) is +3.7 k J/tool for phthalate and +5.0 kJ/mol for salicylate. Because the effect of surface coverage on adsorption enthalpy is unlikely to be greater than its effect on the free energy, no strong dependence of AH on surface coverage (over this concentration range) would be expected.
Compar&on of complex formation in solution and at the mineral-water interface Complex formation in solution is a ligandexchange reaction in which water molecules (or hydroxide ions) are displaced from the inner coordination sphere of the metal. The release of coordinated waters from the relatively constrained metal complex results in a significant increase in entropy; this effect is particularly important for compact, highly charged cations such as A13÷. The formation of the 1:1 aluminum-salicylate complex by the ligand-exchange reaction A1OH 2+ + L 2- ~- AlL + + O H (L = salicylate) proceeds because of a favorable
The adsorption of organic ligands on the surface of aluminum oxide (at pH 8) involves a surface ligandexchange reaction in which a surface hydroxyl group is displaced by the sorbing ligand. Both mono- and bidentate surface complexes may be formed (depending on the structure of the ligand). Adsorption of (partially or fully) protonated ligands results in a neutralization reaction between protons displaced from the ligand and hydroxide ions displaced from the surface (ScHINDLERand STUMM,1987; WESTALL, 1987). The series of organic ligands studied here includes both monodentate ligands (benzoate, the chlorophenols and the aliphatic acids) and bidentate ligands (salicylate and phthalate) some of which are fully deprotonated at the experimental pH (benzoate, phthalate and the aliphatic acids) and others which are partially or fully protonated (salicylate and the chlorophenols). Adsorption isotherms for the aromatic acids exhibit saturation at high dissolved ligand concentrations whereas no saturation was observed (over the concentrations range examined) for the chlorophenols. Small, though measurable, amounts of heat were evolved on the reaction of propionate, the aromatic acids and the chlorophenols with the &A1203 surface; no satisfactory measurements could be made for the fatty acids caprylate and laurate. Overall, reactions of the aromatic acids, particularly salicylate, resulted in the most heat evolution. The calculation of reaction enthalpies from the measured heat evolved was subject to considerable uncertainty because of the necessity of estimating adsorbed ligand concentrations from adsorption isotherms. None the less, examination of the reaction enthalpies may
Adsorption of organic ligands on aluminium oxide
137
Table 8. Enthalpies and free energies of complexation reactions in solution* Reaction
AH (kJ/mol)
AI3+ + HL = AlL + + H + H + + L2- = HLA1OH2-" = AI3+ + OHA1OH2+ + L2- = AlL ÷ + OH-
32.6 -36 +6.1 +2.7
log Kt
AG° (kJ/mol)
0.183 13.74 -9.01 4.91
-1.05 -78.43 51.43 -28.05
* L = salicylate; references: CHR1STENSENet al. (1976); CHRISTENSENand IzA-rr (1983); MACHESKYand JACOaS(1991a); MARTELLand SMITH(1977); t Stability constants for reactions as written.
provide some insight into processes occurring at the mineral surface. The overall reactions of the sorbing ligands with the 6-A1203 surface were exothermic for all ligands (i.e. other than the fatty acids). The relatively small values of AHobs, however, suggest that favorable entropy changes may be important (perhaps even predominant) in the formation of surface complexes. For reaction of protonated ligands (salicylate and the chlorophenols), the favorable enthalpies of reaction (i.e. AHobs) may be due largely to the exothermic proton transfer reaction between the phenolic hydroxy group and hydroxide ions displaced from the oxide surface. If the enthalpic contributions of such reactions are accounted for, the enthalpies of the surface ligand-exchange reaction (AHcorr), that is for the reaction -=AIOH + L"- ~- =-ALL~'-1)- + O H (n = 1 or 2) may be compared for the various ligands. The values of AHcorr appear to be related to the chelating functionality of the ligand; the surface ligand-exchange reaction is more exothermic for phthalate than for benzoate or salicylate and endothermic for the chlorophenols. It should be noted, however, that the surface ligand-exchange reactions would be less exothermic than the values of AHcorr would indicate if other (exothermic) reactions, particularly surface protonation, had contributed to the measured enthalpies. The calculation of AHco~r allows some estimation of the contribution of the surface ligand--exchange reaction to the overall adsorption enthalpy. It is the enthalpy of the surface ligand-exchange reaction that should be most directly related to the structure and bond strength of the surface complex. A pronounced influence of the structure of surface complexes on surface reactivity toward dissolution has been demonstrated for aluminum oxide (FURRER and STUMM, 1986). The extent of A1 and P release from soil mediated by organic acids has been shown to vary with the specific organic acid used (Fox et al., 1990). Thus the mobilization of A1 and of substances, such as P, that adsorb strongly to aluminum oxides from soils will be influenced not only by the concentration but also by the composition of dissolved organic matter in soil waters.
The relation of surface structure and reactivity has implications for the fate and transport of surfaceassociated organics as well as for mineral weathering. The transport of organic substances, including pollutants such as pesticides, in soils and aquifers will be retarded by their adsorption onto immobile solid surfaces. Hydrolysis, an important mechanism for the degradation of thiophosphate pesticides, has been shown to be catalyzed by oxide surfaces when the structure of the organic ester undergoing hydrolysis allows for formation of a bidentate surface c o m p l e x (TORRENTS and STONE, 1991). Thus both the extent and the nature of organic-surface interactions will affect the persistence of organic compounds in the environment. Acknowledgements--We thank Professor R. Giovanoli
(University of Bern) for sample analysis and Professor I. Grenth6 (Royal Institute of Technology. Stockholm) for helpful discussions at early stages of this work. The participation of P. Benoit was sponsored by the Institut National de la Recherche Agronomique, Thiverval-Grignon, France. This work was supported in part by the Swiss National Science Foundation. Editorial handling: J. G. Catts.
REFERENCES
AMAL R., RAPERJ. A. and WAITET. D. (1992) Effect of fulvic acid adsorption on the aggregation kinetics and structure of hematite particles. J. Colloid Interface Sci. 151,244-257. BALASTRIERIL. S. and CHou T. T. (1987) Selenium adsorption by goethite. Soil Sci. Soc. Am. J. 51, 1145-1151. CHRISTENSENJ. J., HANSENL. D. and IZAYrR. M. (1976) Handbook of Proton Ionization Heats. Wileylnterscience. CHRISTENSEN J. J. and IZATr R. M. (1983) Handbook of Metal Ligand Heats. Marcel Dekker. DAVISJ. A. (1982) Adsorption of natural dissolved organic matter at the oxide/water interface. Geochim. Cosmochim. Acta 46, 2381-2393. DAVIS J. A. and GLOOR R. (1981) Adsorption of dissolved
organics in lake water by aluminum oxide. Effect of molecular weight. Envir. Sci. Tech. 15, 1223-1229. DENOYEL R., DURAND G., LAFUMA F. and AUDEBERT R.
(1989) Adsorption of cationic polyelectrolytes onto montmorillonite and silica: a microcalorimetric study of their conformation. J. Colloid Interface Sci. 139,281-290. DOUGANW. K. and WILSONg. L. (1974) Absorptiometric determination of aluminum in water. Comparison of
138
P. Benoit et al.
some chromogenic reagents and the development of an improved method. Analyst 99,413-430. DZOMBAK D. A. and MOREL F. M. M. (1990) Surface Complexation Modeling: Hydrous Ferric Oxide. WileyInterscience. EATHOUGH D. J., CHRISTENSENJ. J. and IZATr R. M. (1974) Experiments in Thermometric Titrimetry and Titration Calorimetry. Brigham Young University Press. FOKKINK L. G. J., DE KEIZER A. and LYKLEMAJ. (1989) Temperature dependence of the electrical double layer on oxides: Rutile and hematite. J. Colloid Interface Sci. 127, 116-131. FOKKINK L. G. J., DE KEIZER A. and LYKLEMAJ. (1990) Temperature dependence of cadmium adsorption on oxides. J. Colloid Interface Sci. 135, 118-131. Fox T. R., COMERFORD N. B. and MCFEE W. W. (1990) Phosphorous and aluminum release from a spodic horizon mediated by organic acids. Soil Sci. Soc. Am. J. 54, 1763-1767. FURRER G. (1985) Die oberflaechenkontrollierte Aufloesung von Metalloxiden: ein koordinationschemischer Ansatz zur Verwitterungskinetik. Ph.D. Thesis, Swiss Federal Institute of Technology, Zurich. FURRER G. and STUMMW. (1986) The coordination chemistry of weathering: I. Dissolution kinetics of 6-AbeO3 and BeO. Geochim. cosmochim. Acta 50, 1847-1860. HAYES K. F., PAPELISC. and LECKIEJ. O. (1988) Modeling ionic strength effects on anion adsorption at hydrous oxide/solution interfaces. J. Colloid Interface Sci. 125, 717-726. HER1NG J. G. and STurau W. (1991) Fluorescence spectroscopic evidence for surface complex formation at the mineral-water interface: elucidation of the mechanism of ligand-promoted dissolution. Langmuir 7, 1567-1570. HUNTER K. A. and LIss P. S. (1979) The surface charge of suspended particles in estuarine and coastal waters. Nature 282,823-825. JAFVERT C. T. and WOLFE N. U (1987) Degradation of selected halogenated ethanes in anoxic sediment-water systems. Envir. Toxicol. Chem. 6, 827-837. JEKEL M. R. (1986) The stabilization of dispersed mineral particles by adsorption of humic substances. Water Res. 12, 1543-1554. DE KEIZER A., FOKKINK L. G. J. and LYKLEMAJ. (1990) Thermodynamics of proton charge formation on oxides. Microcalorimetry. Colloids Surf. 49,149-163. KUMMERTR. and STUMraW. (1980) The surface complexation of organic acids on hydrous 7-A1203. J. Colloid Interface Sci. 75,373-385. KUNG K.-H.S. and MCBRIDE M. B. (1991) Bonding of chlorophenols on iron and aluminum oxides. Envir. Sci. Tech. 25,702-709. LIANG L. and MORGANJ. J. (1990) Chemical aspects of iron oxide coagulation in water: laboratory studies and implications for natural systems. Aquatic Sci. 52, 32-55. MACALADV D. L. and WOLFE N. L. (1984) Abiotic hydrolysis of sorbed pesticides. In Treatment and Disposal of Pesticide Wastes (eds R. F. KRUEGER and J. N. SEIaER), pp. 221--244, ACS Syrup. Ser. 259. MACHESKYM. L. and ANDERSON M. A. (1986) Calorimetric acid-base titrations of aqueous goethite and futile suspensions. Langmuir 2, 582-587. MACHESKY M. L., BISCHOFF B. L. and ANDERSON M. A. (1989) Calorimetric investigation of anion adsorption onto goethite. Envir. Sci. Tech. 23, 580-587. MACnESRV M. L. and JACOBSP. F. (1991a) Titration calorimetry of aqueous alumina suspensions Part I. Results and comparisons with similar studies. Colloids Surf. 53,297314. MACHESKYM. L. and JACOBSP. F. (1991b) Titration calorimetry of aqueous alumina suspensions Part II. Discussion of enthalpy changes with pH and ionic strength. Colloids Surf. 53,315-328.
MARTELL A. E. and SMITh R. M. (1977) Critical Stability Constants, Vol. 3. Plenum Press. MCBRtDE M. B. (1980) Adsorption of fatty acid spin probes on amorphous alumina. J. Colloid Interface Sci. 76,393398. MCBRIDE M. B. (1982) Organic anion adsorption on aluminum hydroxides: Spin probe studies. Clay Clay Minerals 30,438-444. NOEL L. A. (1987) Adsorption calorimetry of surfactant interaction with minerals. Colloids Surf. 26, 43-54. OCHS M., BRUNNER I., STUMMW. and t~osowC B. (1993) Effects of root exudates and humic substances on weathering kinetics. Water Air Soil Pollut. (in press). PAPEEIS C., HAYES K. F. and LECKIE J. O. (1988) HYDRAQL: a program for the computation of chemical equilibrium composition of aqueous batch systems including surface complexation modeling of ion adsorption at the oxide/solution interface. Tech. Rep. No. 306. Dept. of Civil Engineering. Stanford Univ., Stanford, California. PARTYKAS., LINDHEIMERM., ZA1NIS., KEH E. and BRUN B. (1986) Improved calorimetric method to investigate adsorption processes from solution onto solid surfaces. Langmuir 2, 101-105. PARTYKA S., RUDZUNKA W., BRUN B. and CLINT J. H. (1989) Calorimetric studies of adsorption of anionic surfactants onto alumina. Langmuir, 5,297-304. SANDERS P. and WOLF N. L. (1985) Reduction of nitroaromatic compounds in anaerobic, sterile sediments. Presented at the National Meeting of the Am. Chem. Soc., Chicago, Illinois. SCHINDLERP. W. and STUMMW. (1987) The surface chemistry of oxides, hydroxides, and oxide minerals. In Aquatic Surface Chemistry (ed. W. STUMM), pp. 83--110. WileyInterscience. STONE A. T. (1987) Reductive dissolution of manganese (Ill/IV) oxides by substituted phenols. Envir. Sci. Tech. 21,979-988. STUMM W., KUMMERT R. and SIGG L. (1980) A ligand exchange model for the adsorption of inorganic and organic ligands at hydrous oxide interfaces, Croatica Chem. Acta 53,291-312. STUMMW. and WIELANDE. (1990) Dissolution of oxide and silicate minerals: Rates depend on surface speciation. In Aquatic Chemical Kinetics (ed. W. STUMM),pp. 367-400. Wiley-lnterscience. TEJEDOR-TEJEDOR M. I., YOST E. C. and ANDERSONM. A. (1990) Characterization of benzoic and phenolic complexes at the goethite/aqueous solution interface using cylindrical internal reflection Fourier transform infrared spectroscopy. Part I. Methodology. Langmuir 6, 979987. TIPPING E. (1981) The adsorption of aquatic humic substances by iron oxides. Geochim. Cosmochim. Acta 45, 191-199. TIPPING E. (1986) Some aspects of the interactions between particulate oxides and aquatic humic substances. Mar. Chem. 18,161-169. TIPPING E. (1990) Interactions of organic acids with inorganic and organic surfaces. In Organic Acids in Aquatic Ecosystems (ed. E. M. PERDUEand E. T. (GJESSING),pp. 209--221. John Wiley and Sons. TORRENTS A. and STONE A. T. (1991) Hydrolysis of phenyl picolinate at the mineral/water interface. Envir. Sci. Tech. 25,143-149. ULRICH H.-J. and STONE A. T. (1989) Oxidation of chlorophenols adsorbed to manganese oxide surfaces. Envir. Sci. Tech. 23,421-428. ULRICH H.-J., STUMMW. and (~osovI~ B. (1988) Adsorption of aliphatic fatty acids on aquatic interfaces. Comparison between two model surfaces: the mercury electrode and ~-A120 3 colloids. Envir. Sci. Tech. 22, 37-41. WAKAMATSUT. and FEURSTENAUD. W. (1968) The effect of
Adsorption of organic ligands on aluminium oxide hydrocarbon chain length on the adsorption of surfactants at the solid/water interface. In Adsorption from Aqueous Solution (ed. W. I. WEBER and E. MATEJWlC), pp. 161--172. ACS Adv. Chem. Ser. 79. WEBER E. J. and WOLFE N. L. (1978) Kinetic studies of the reduction of aromatic azo compounds in anaerobic sediment/water systems. Envir. Toxicol. Chem. 6, 911919. WESTALL J. C. (1987) Adsorption mechanisms in aquatic surface chemistry. In Aquatic Surface Chemistry (ed. W. STUMM)pp. 3--32. Wiley-lnterscience. WESXALLJ. C., ZACHARYJ. L. and MORELF. M. M. (1976) MINEQL: a computer program for the calculation of the chemical equilibrium composition of aqueous systems, R. M. Parsons Laboratory Tech. Note No. 18, Massachusetts Institute of Technology, Cambridge.
139
WIERERK. A. and DOBIASB. (1988) Exchange enthalpies of H + and O H - adsorption on minerals with different characters of potential-determining ions. J. Colloid Interface Sci. 122, 171-177. YAPS. N., M1SHRAR. K., RAGHAVENS. and FUERSTENAUD. W. (1981) The adsorption of oleic acid on goethite. In Adsorption from Aqueous Solutions (ed. P. H. TEWARI), pp. 119--142. Plenum Press. ZELTNER W. A., YOST E. C., MACHESKYM. L., TEJEDORTEJEDOR M. I. and ANDERSONM. A. (1986) Characterization of anion binding on goethite using titration calorimetry and cylindrical internal reflection-Fourier transform infrared spectroscopy. In Geochemical Processes at Mineral Surfaces (ed. J. A. DAVIESand K. F. HAYES),pp. 142--161. ACS Symp. Ser. 323.