Environ. Sci. Technol. 1993, 27, 895-909
Adsorption of Organic Compounds Possessing Ligand Donor Groups at the Oxide/Water Interface Alan T. Stone,. Alba Torrents, Jean Smolen, Dharanlja Vasudevan, and James Hadley
Department of Geography and Environmental Engineering, G. W. C. Whiting School of Engineering, Johns Hopkins University, Baltimore, Maryland 212 18
A diffuse-layer model is used to demonstrate that the effects of pH and ionic strength on organic ligand adsorption are a direct consequence of (i) stoichiometry of surface complex formation; (ii) mass balance equations for ligands, protons, and surface sites; and (iii) PoissonBoltzmann terms which relate ion concentrations at the oxidelwater interface to concentrations in bulk solution. Experiments examining the adsorption of 2,4-dinitrophenol, 2-pyridinemethanol, and several substituted 2-aminophenols onto TiO2, A1203, and FeOOH are used to illustrate the application of this approach. Once oxide surface protonation equilibria are properly accounted for, ligand adsorption can be modeled effectively in most instances through the choice of a single reaction stoichiometry and a single fitting parameter, the surface complex formation constant. Introduction Planned and inadvertent release of organic chemicals into soils, sediments, and aquifers brings them into contact with particulate natural organic matter (PNOM) and mineral surfaces. As contaminated waters infiltrate porous media, several processes are expected to attenuate the movement of organic chemicals: phase partitioning into PNOM (1, 2), ion exchange (3), and adsorption onto mineral surfaces ( 4 ) . The present work seeks to provide the scientific basis for evaluating the importance of adsorption onto mineral surfaces, particularly in subsurface media where the PNOM content is low. Even in situations where adsorption onto mineral surfaces is not sufficiently strong to attentuate movement, it can still exert a profound effect on pathways and rates of organic pollutant degradation. Because of solubility limitations, the oxidation of phenols (5, 6), anilines (7), and mercaptans (8,9) by Fe(II1)and Mn(II1,IV)necessarily involves adsorption onto mineral surfaces. In some instances, the hydrolysis of carboxylic acid esters and phosphorothioate pesticides can be accelerated via adsorption onto Al(II1)-,Fe(III)-,and Ti(1V)-bearingmineral surfaces ( 1 0 , I I ) . Thus, new information concerning the adsorption of organic compounds can improve our understanding of surface chemical reactions. We are seeking ways of predicting the extent of organic chemical adsorption onto mineral surfaces based upon the physical and chemical characteristics of the adsorbing organic chemical, the mineral surface, and the supporting aqueous medium. As part of this effort, the adsorption of organic chemicals containing one or more ionic or ionizable functional groups onto A1203, TiO2, and FeOOH has been examined in our laboratory. The approach taken here is a logical extension of earlier efforts by Kummert and Stumm (121, Davis and Leckie (13),Balistrieri and Murray (141, Zachara et al. (41, and others.
* To whom all correspondence should be addressed. 0013-936X/93/0927-0895$04.00/0
0 1993 Amerlcan Chemical Society
Adsorption depends upon physical and chemical interactions taking place at the oxidelwater interface. Several models of varying complexity are available for accounting for these interactions (15), ranging from relatively simple models with few fitting parameters such as the diffuse-layer model to intricate models involving numerous fitting parameters such as the triple-layer model. Our discussion focuses upon identifying system parameters that have the greatest effect on ligand adsorption and deciding on the level of complexity required in a model in order to represent experimentally observed phenomena.
Theory Our system consists of an oxide surface, an adsorbing organic compound, the supporting electrolyte, and the solvent water. In most studies, the extent of adsorption is determined by measuring loss of test organic compound from solution and applying a mass balance equation to calculate the amount of organic compound adsorbed. “Apparent” adsorption constants derived from such measurements reflect the combined contributions from all physical and chemical forces influencing adsorption. These forces can be divided into two general categories. “Specific” adsorption involves physical and chemical interaction of the adsorbing chemical with individual sites on the oxide surface; the chemical nature of the sites (e.g., the identity of surface-bound metal atoms) influences the extent of adsorption. “Nonspecific” adsorption, in contrast, does not depend upon the chemical nature of individual surface sites. Instead, the aggregate characteristics of the oxide/water interface, such as the surface charge density, determine the magnitude of the nonspecific contribution toward adsorption. As our knowledge of these forces is improved, we may elect to treat each contribution explicitly. Our knowledge of long-range and short-range electrostatic forces and their effect on adsorption, for example, is relatively advanced. The next section discusses how apparent adsorption constants can be differentiated into two terms, one representing long-range electrostatic interactions, and the other representing all remaining physical and chemical interactions. Long-Range Electrostatic Interactions. Although solvent-solvent, solvent-organic, solvent-electrolyte, and electrolyte-organic interactive forces may change in the vicinity of the oxidelwater interface, the magnitude of these changes is difficult to evaluate. For low molecular weight organic chemicals, most researchers assume that these changes are small and that nonspecific adsorption arises primarily from the action of long-range electrostatic forces on counterions in the vicinity of charged interfaces. In the absence of any specific chemical interaction, ions freely moving within the diffuse layer attain an equilibrium distribution dictated by the requirement that the electrochemical potential (ELI) is constant throughout the system. For an ion of charge Zi, cii at distance x: from the Environ. Scl. Technol., Vol. 27, No. 5, 1993 895
surface can be divided into a chemical term (pi(x))and an electrostatic term (26): ai(x)
= ~ i ( x+ ) ziF$,
calculate the fraction adsorbed by long-range electrostatic interaction:
(1)
F is the Faraday constant, and $., is the electrical potential at distance x from the surface. The magnitude of $, is determined by the electrical potential at the plane of closest approach and by the ion density in solution. We will employ the diffuse-layer model here (see ref 15),which equates with $d, the diffuse-layer potential. The chemical potential pi(x) of species i is related to its activity ({i))and concentration ([il), according to the following equation: p i ( x ) = poi
+ R T In {i)= poi + RT In yi(x)[il,
(2)
yi(x) is the activity coefficient at distance x from the
surface. Combining eqs 1 and 2 reveals that the activity of species i at distance x from the surface is related to the activity in bulk solution times a Poisson-Boltzmann factor (see ref 17): {i), = {i)bulkexp(-ziFICX/RT)
(3)
It will be assumed that the activity coefficient (yi) for species i is independent of distance from the charged surface. With this constraint, concentrations at distance x from the surface are equal to the bulk concentration times the Poisson-Boltzmann factor. We calculate the buildup of species i in the diffuse layer by noting that the total moles in solution (ni) can be found by integrating eq 3 over the entire volume of solution:
where S is the oxide surface area (cm2), and V is the suspension volume (cm3). Based upon this result, the amount nonspecifically adsorbed (the excess in the diffuse layer) can be found by mass balance: fraction adsorbed =
(ni/V)- [iIbulk ni/ V (S/v)rnv's exp(-ziF$,/RT) dx - 1
We will restrict ourselves to the nonspecific adsorption of a monovalent anion (L-). In a 1:l electrolyte solution, #, can be calculated from $, and x using the following equation (18): exp(F$,/RT) = [exp(F$,/2RT) [exp(F$,/2RT)
(
+ 11 + [exp(F$JZRT) - 11 exp(-ax) * + 11 - [exp(F$J2RT) - 11 exp(-Kx)
1
(6)
The characteristic thickness of the diffuse layer 1 / (ref ~ 19, p 617) is a function of the ionic strength p , as given by the following equation: K
= 3.57 x 10%'; (in cm-l)
(7)
Inserting the solution given below (Z. Z. Zhang and Donald Sparks, personal communication) into eq 5 allows us to 896
Environ. Sci. Technol., Vol. 27, No. 5, 1993
In this equation, I = [exp(F$,/2RT) + 11 and J equals [exp(F$,/BRT) - 11. Note that INTEG is equal to unity when the surface bears no net charge. Thus, the extent of nonspecific adsorption arising from long-rangeelectrostatic interactions can be calculated once the electrical potential at the oxide surface is known. In order to calculate the extent of nonspecific adsorption, specific chemical interactions which determine surface charge must be discussed. Near-Range Physical and Chemical Interactions. As has already been mentioned, specific adsorption involves interaction with individual sites on the oxide surface. These interactions can be accounted for in two ways. The first approach involves using activity coefficients to relate the electrochemical activity of a species at the oxide/water interface to its electrochemical activity in bulk solution. This approach is most suited for situations where the making or breaking of covalent bonds is not taking place and where an established physical or chemical relationship can be used to estimate activity coefficients. Long-range electrostatic forces have already been accounted for using the Poisson-Boltzmann term (eq 3). An additional correction accounting for short-range electrostatic forces may be appropriate when the proximity of the adsorbate to the surface is closer than the site-to-site charge density. To accomplish this, it would be necessary to define an effective radius for the charge located on the surface site and for the adsorbing ion. In the second approach, a new chemical species is postulated, and an equilibrium constant is defined that relates the activity of the new species to the activities of established species. This approach is the appropriate one when the making or breaking of covalent bonds is taking place. It can also be used in situations where the physical or chemical forces involved in adsorption are not known with certainty or are difficult to quantify. Because equilibrium constants are mass action expressions, it is important that reaction stoichiometry be critically evaluated. When the postulated stoichiometry is correct, it should be possible to predict changes in the concentration of the surface species as the concentrations of other constituent chemicals are altered. Surface Complex Equilibria and Assignment of Stoichiometry. We begin our modeling effort by making the following assumptions: (i) all surface sites are chemically identical, (ii) all species adsorb at the same plane, (iii) dissolved species near the oxide/water interface have activity coefficients identical to those of species in bulk solution, and (iv)adsorbed specieshave activity coefficients equal to 1.0. Physical and chemical characteristics of the adsorbate ligands and the oxide surfaces of interest here are included in Tables I and 11. The diffuse-layer model will be used because it is the simplest model that can still account for the selected level of complexity. The computer program MINEQL (23),which formalizes the solution of mass action and mass balance equations, has been modified by various groups to treat surface charge, surface potential, and other aspects of the electrical double layer as well.
Table I. Acidity Constants and Surface Complex Formation Constants for the Ligands Examined in This Study adsorption conditions log aK7,aintr compound PKa 2,4-dinitrophenol(2,4-DNP)
4.11 (20)
2,6-dinitrophenol(2,6-DNP) 4-methyl-2-aminophenol(4-Methyl-2-AP) 2-aminophenol(2-AP) 3-amino-2-naphthol (3-amino-2-NAP) 4-nitro-2-aminopheno1(4-nitro-2-AP) 2-pyridinemethanol(2-PM)
3.70 (20) 4.3,10.0a 4.7,9.9 (21) 3.9, 9.2a 3.1,7.6” 4.86,13.9 (22)
7.1
10 g/L Ti02 (low ionic strength)
6.5 12.2
10 g/L FeOOH (low ionic strength) 10 g/L A1203 (low ionic strength) 10 g/L Ti02 (low ionic strength) 10 g/L Ti02 (0.10 M NaCl)
11.9 11.6 10.9 15.7
10 g/L Ti02 (0.10 M NaCl) 10 g/L Ti02 (0.10 M NaCl) 10 g/L Ti02 (0.05 M NaC1) 7 g/L TiOz (0.11 M NaCl)
6.7
This study. Table 11. Properties of Synthetic Oxides Employed in Adsorption Experimentsa oxide A1203 (Type C) FeOOH (Goethite) Ti02 (Type P25)
site density, sites/nm2
S, m2/g pH,,,
92.8 30.8 55.0
8.9 107,7 10-10.2 7.gb 1O6.O * 6.5 103.9 10-8.7
2-4 6-7b 3
a sK3intr and aKqintrare defined by reactions 3 and 4 in Table 111. Values were determined for another FeOOH suspension, synthesized according to the same procedure (from ref 10). ~
~~
Here we will use the revised program HYDRAQL (24). Details concerning the operation of these programs and inclusion of the diffuse-layer model are available in the excellent review by Westall and Hohl (15). A comprehensive treatment of the adsorption of inorganic anions and cations onto oxides is available in Dzombak and Morel (25). Table I11 lists chemical reactions and corresponding mass action expressionsrequired by the diffuse-layermodel in order to model adsorption of the generic ligand L-. For monoprotic ligands such as 2,4-DNP and 2,6-DNP,reaction 1 (HLO/L-) is sufficient for describing the protonation speciation in solution. For diprotic ligands such as substituted aminophenok, reactions 1and 2 (H2L+/HLo/L-) are both required. Reactions 3-10 account for equilibria among surface species. In this work we will use a conventional “threestate model” (26), which has been found to realistically repwent surface protonation as a function of pH during acid/base titrations. Application of the three-state model to outer-sphere complex formation (reactions 5 and 7) is conceptually straightforward; the inner coordination sphere of the surface-bound metal atoms at each site remain intact. When inner-sphere surface complex formation takes place, the protonation level of neutral, unoccupied surface sites and the proton stoichiometry of the adsorption process become important issues. The notation most commonly used for the neutral, unoccupied surface site (>MOH) is a reasonable one for denoting surface reactions where only one coordinative position of the surface-bound metal is involved. Consider reaction 8a as it applies to the adsorption of a monodentate ligand: a proton is consumed in the conversion of OH- into HzO, which opens up a vacancy in the inner coordination sphere of the surfacebound metal that can be then occupied by an incoming ligand. This notation adequately represents the adsorption of monodentate ligands and has been extended in some instances to describe bridging by bidentate ligands across two surface sites (e.g., (>M)zL+, as used by Sigg and Stumm (27)).
When two or more coordinative positions of the surfacebound metal are involved in surface complex formation, however, the conventional notation is inadequate. In order to simplify comparisons among alternative stoichiometries, all surface complex equilibrium reactions are written beginning with the unoccupied surface site in its neutral form. On this basis, the central issue is whether zero, one, or two protons must be adsorbed in order to provide two inner-sphere coordinative positions for an incoming bidentate ligand. Stated differently, the issue is whether ligand adsorption causes an increase, a decrease, or no change in the charge of the surface site. Three alternative protonation stoichiometries for the adsorption of bidentate ligands are represented in reactions 8b, 9, and 10 in Table 111. If the neutral, unoccupied site is best represented as >M(OH)(OHz), then consumption of a single proton is required for bidentate ligand adsorption to take place. If the stoichiometry >M(OH)2 is the most appropriate representation, then two protons must be consumed upon entry of the bidentate ligand, and the surface site charge increases by one unit. If >M(OH& is the most appropriate representation, then no protons are consumed during the adsorption process; the surface site charge decreases by one unit. The assignment of protonation stoichiometry can be considered a fitting exercise: the choice can be made based upon the ability of the model to represent the experimental data. In our diffuse-layer model, reactions 7,8a, and 8b in Table I11 are mathematically equivalent: consumption of one proton must accompany ligand adsorption. It is easy to recognize that alternative protonation stoichiometries represented by reactions 9 and 10 affect the pH dependence of the adsorption reaction as predicted by the model. It is also true that these alternative stoichiometries alter ionic strength effects. At fixed surface charge, long-range electrostatic interactions are strongest at low ionic strength. On positively charged oxide surfaces, long-rangeelectrostatic interactions impede the formation of positive surface complex >ML+ (reaction 9), enhance formation of negative surface complex >ML- (reaction lo), and leave formation of neutral surface complex >MLo (reactions 7,8a, and 8b) unperturbed. These issues will be discussed in greater detail in later sections. Other lines of evidence can help resolve the proton stoichiometry question. Hiemstra et al. (28,29) have used bulk structural and coordinative characteristics of oxide minerals to estimate the protonation behavior of surface sites. Ti(1V) atoms within the mineral rutile (TiO2) are octahedrally coordinated; in order for the bulk solid to be electrically neutral, an average bond valence of 2/3 can be assigned to each Ti-0 bond (28). We now shift our attention to a Ti atom on the oxide surface that has four Envlron. Scl. Technol., Vol. 27, No. 5, 1993
897
Table 111. Reactions Pertinent to Diffuse-Layer Model and Corresponding Mass Action Equations
L- + H+ = HLO L- + 2H+ = H2L+ >MOH + H+ = >MOHz+ >MOH = >MO- + H+ >MOH + H+ + C1- = >MOHz+Cl>MOH + Na+ = >MO-Na+ + H+ >MOH + H++ L- = >MOHz+L>MOH + H+ + L- = >ML + 2H20 >M(OH)(OH,) + H++ L- = >ML + HzO >M(OH)? + 2H+ + L- = >ML+ + 2H20 >M(OH2)2 + L- = >ML- + 2Hz0
bonds to the oxide lattice and two coordinative positions vacant. Summing the charges contributed by all six innersphere ligands, a charge of +1/3 is assigned to surface site >Ti(OH)(OHz). Similar treatment of surface sites >Ti(0H)z and >Ti(OHz)2 yields charges of -2/3 and +4/3, respectively. Thus, an electrically neutral site with two exchangeable coordinative positions comes closest to the stoichiometry represented by reaction 8b in Table 111. Mass Action Equations. Mass action expressions listed in Table I11are “activity scale”equilibrium constants expressed in terms of solute activities. By this approach, nonideality arising from ionic interactions in solution has been accounted for; equilibrium constants involving only solute species (aK1and a K ~are ) true constants, possessing values independent of ionic strength. Additionally, equilibrium constants involving surface species have been expressed as “intrinsic” constants, denoted by the superscript “intr”, signifying that mass action is expressed in terms of solute activity at the plane of closest approach to the surface (&). Intrinsic surface constants are also true constants, possessing values independent of ionic strength. Alternatively, equilibrium constants can also be expressed as “apparent” constants, denoted by the superscript “app”, signifying that mass action is expressed in term of solute activity in bulk solution. Apparent surface constants are no longer true constants: each charged solute included in the reaction stoichiometry carries along an electrostatic correction term relating its activity at the plane of closest approach to its activity in bulk solution; each Poisson-Boltzmann electrostatic correction term changes value whenever the properties of the electrical double layer are changed. In order to solve the mass balance equations, all dissolved speciesactivities must be converted to concentrations using 898
Envlron. Sci. Technol., Vol. 27, No. 5, 1993
the Davies equation (ref 19, p 135),and concentrations at the plane of closest approach must be related to bulk solution concentrations using the Poisson-Boltzmann correction term. Transformation of the mass action equations in this manner is presented in Table IV. In keeping with the assumption that adsorbed species have activity coefficients equal to 1.0, activity notation (e.g., (>MOHz+J)has also been dropped. Concentrations of surface sites (e.g., [>MOH2+]) are given in units of moles of sites/liter. Using the assumption that all swface sites are chemically identical, we will now represent any neutral, unoccupied surface site as simply “>MOH”. Table V lists concentrations of all species pertinent to the diffuse-layer model calculation in terms of three component concentrations: [H+]bulk,[L-lbulk, and [>MOH]. These equations will be used in the next section to formulate mass balance equations. Of the 10 species listed, nine include an activity coefficient correction. When no net change in surface site charge accompanies a surface chemical reaction (e.g., reactions 7 and 8) the Poisson-Boltzmann terms cancel out. For these reactions, changes in electrolyte concentration and accompanying changes in long-range electrostatic forces have little direct effect on the extent of surface complex formation. Reactions where a net change in surface charge does take place (e.g., reactions 9 and 10) are much more sensitive to changes in ionic strength. Proton concentrations and Poisson-Boltzmann correction terms are found in many competing terms within each mass balance equation. For this reason, assessing indirect effects of changes in pH and electrolyte concentration can be quite complex. Reactions 5-7 represent outer-sphere complex formation. If the triple-layer model is used in place of the diffuselayer model, adsorption of H+ would be assigned to the
Table IV. Mass Action Equations Expressed in Terms of Bulk Solution Concentrations, Using Activity Corrections and Poisson-Boltzmann Terms
'K, =
[HLol -- -cKl [L~lbu]kY.[H+lbu]kY+ Y-Y+
(1)
Table V. Concentrations of Pertinent Species in Terms of [ H + l b ~ i k , [ L - l b u i k , and [>MOW
[HLOI = a ~ ~ ~ - ~ - [ ~ ~ l b ~ l k [ ~ ~ l b u l k (1) [HzL+I = "K~Y-Y+~[L~lbulk[H+]~bulk (2) [>MOH2+] = aK3"'trY+ e x ~ ~ - F $ ~ / ~ T ) ~ ~ +(3)l b ~ ~ k ~ ~ ~ ~ ~ ~ aK4'ntrexp(+F&,/RT) [>MOHl [>MO-] = (4) Y+ LH+] bulk [>MOH,+CI-l = a ~ ~ ' n t r ~ + ~ - ~ ~ ~ (5) ~ b ~ ~ k ~ ~ ~ ~ ~ b u [Na+lbulk[>MOHl [>MO-Na+] = aKglntr (6) [H+lbu,k [>MOHz+L-] = aK7'ntrY+Y-[H+]buik[L-]buik[>MOH] (7) [>MLl = a K ~ i " t ' ~ + ~ - tbulk[L-lbu~k[>MOHl H+l (8) [>ML+l = (9) aKgintrY+2Y-exP(-F$o/RT) [H+12b~~k[L~lbu~k[>MOHl PML-I = aKln'"fry-eXp(+F$~/RT)[L-lbulk[>MOH1 (10) 1 2 3 4
innermost plane, while adsorption of C1-, Na+,and L- would be assigned to a second plane, a short distance away from the surface. In this case, the Poisson-Boltzmann correction terms would not completely cancel out in reactions 5-7, and the effect of changes in ionic strength on complex formation would be stronger. In practice, assignment of different adsorption planes is an artificial one. In the absence of specific evidence concerning molecular distances at the surface, the different adsorption planes and associated changes in Poisson-Boltzmann correction terms become additional fitting parameters. Calculating the Extent of Adsorption from Mass Balance Equations. Each modeling problem requires that different chemical species be evoked. In order to examine how long-range electrostatic interactions influence adsorption, illustrative cases involving the species listed in Table I11 will be examined. A 10 g/L goethite
~
k
5
~
6
~
7
~
~
8
~
9
~
1
0
1
1
PH
Flgure 1. Calculated FeOOH surface charge (expressed as moleequivalent charge divided by &) as a function of pH for four concentrations of indifferent electrolyte. Model conditions: 10 g/L FeOOH, yielding a surface loading of 501 m2/L, and & equal to 5.41 x 10-3 M.
(FeOOH) suspension exhibiting the physical and chemical characteristics presented in Table I1 provides a representative surface for our modeling efforts. For this set of calculations, it will be assumed that specific adsorption of electrolyte ions does not take place and that the moles of ligand added to the system is so low that ligand adsorption does not significantly perturb the surface charge. Thus, charging of the oxide surface arises solely from reactions 3 and 4. Figure 1presents surface charge as a function of pH in FeOOH suspensions containing increasing amounts of Envlron. Sci. Technol.. Vol. 27,
No. 5, 1993 899
,
l.O,,
,
,
I
~,
,
,
,
I
,
I
,
,
,
,
,
,
,
Ionic S t r e n g t h ( m a l e - e q u i v / L )
7-1:: 1
Ionic S t r e n g t h (mole-equiv./L)
0.2
1
\
0 0 10-1,
t -0.21 1
'
1
2
'
1
3
'
4
1
'
5
" " '
1
6
7
8
1
9
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1
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0
1
1 1
Nonspecific adsorption of
(A)
1
i '
I
2
'
1
3
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I
" '
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4
5
6
7
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8
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1 0 1 1
monovalent anion L- and (B) weak acid HLa/L- (paK, 4.11) onto the model goethite surface.
indifferent electrolyte. Surface charge is expressed in units of mole-equivalent charge divided by the total number of surface sites. Intrinsic surface acid/base constants (aK3intr and aK4intr) are derived from acid/base titration. A positive surface charge develops on the oxide surface as the pH drops below the pH,,,. The magnitude of this surface charge increases dramatically as the ionic strength is increased. Two related phenomena are responsible. As the ionic strength is increased, the density of counterions per fluid volume is greater, and therefore the developing surface charge is more readily counterbalanced (the diffuse layer thickness, related to 1 / ~is, decreased). Equally important, the surface potential & decreases as the ionic strength is increased, lessening electrostatic repulsion as the proton approaches the positive oxide surface, as represented in the Poisson-Boltzmann equation (eq 3). Case A, presented in Table VI, involves the nonspecific adsorption of a monovalent anion L-. Comparing the mass balance equation for ligand with eq 5 yields: fraction adsorbed = (INTEG - l)/INTEG
(9)
Results are presented in Figure 2A as a function of pH. Because H+(aq)and OH-(aq) are included in calculations of ionic strength, a suspension bearing a fixed ionic strength mol equiv/L cannot have a pH below 4.0 or of 1.0 X above 10.0. Even though the surface charge is low when the ionic strength is low, the extent of nonspecific adsorption is high. For the same reason that approach of the co-ion H+(aq) is electrostatically unfavorable at low ionic strength, approach of counterions, including the anionic adsorbate, is strongly favorable. One interesting consequence of this calculation is that at ionic strengths 900
-0.21
PH
PH
Flgure 2.
10-2
t
Envlron. Sci. Technol., Vol. 27, No. 5, 1993
below 1.0 mM a substantial fraction of the total counterions in suspension is found in the vicinity of the oxidelwater interface. Case B represents the nonspecific adsorption of a ligand exhibiting acid/base speciation (HLOIL-). A paKaof 4.11 has been selected, the value exhibited by 2,4-dinitrophenol. The concentration (and activity) of the neutral species HLO in the interfacial region and in bulk solution is the same, while L- becomes accumulated near the positive oxide surface. In order to calculate the fraction adsorbed, eq 5 must be modified:
= (INTEG - l)/(M
+ INTEG)
(10)
where M represents aKly+y-[H+]bulk. Results of this calculation are presented in Figure 2B. The extent of nonspecific adsorption initially increases as the pH is decreased, attaining a maximum near the pKa of the adsorbate. A t pH values below the pKa, the extent of nonspecific adsorption diminishes substantially, as the speciation shifts from L- to HLO. Protonation narrows the pH range where nonspecific adsorption can take place and dampens the extent of nonspecific adsorption in the vicinity of the pKa. For compounds possessing higher pKa values, the contribution of nonspecific forces toward adsorption is even less. In Case C, both nonspecific and specific adsorption of ligand HLOIL- take place; reaction 8 in Table I11 is used to represent the stoichiometry of the surface complex.
Table VII. Model Calculation of Extent of Nonspecific Adsorption of an Anion with No Acid/Base Properties (L-) and a Weak Base (HLa/L-) Possessing a pK, of 4.11*
suspension 5.4 gIL A1203 10.0 g/L FeOOH 9.1 giL Ti02
% adsorbed surface charge species HLOILpH,,, density, C/m2 species L(pK, 4.11)
8.9 7.9 6.5
0.18 0.070 0.011
45.3 24.4 4.2
27.0 12.6 1.9
a The oxide surface area (501 m*/L),pH (4.0), and ionic strength (1.0 mM) are fixed.
The fraction adsorbed can be calculated using the following equation:
= (INTEG - 1 + N ) / ( M + INTEG + N ) (11) where N represents aK~intr[H+l~,lk[>MOHl. Clearly, the magnitude of the surface equilibrium constant determines the importance of specific versus nonspecific adsorption in this equation. When aKsintris sufficiently small, eq 11 reduces to eq 10. As we have seen, the importance of nonspecific adsorption on the FeOOH surface decreases substantially as the ionic strength is increased and as the pK, of the ligand HLO/L- is increased. We will now examine how differences in the protonation behavior of FeOOH, TiOz, and A1203 affect the importance of nonspecific adsorption. Oxide loading (g/L) has been adjusted so that the surface area loading in all three suspensions is the same (501 m2/L). The pH has been fixed at 4.0, and the ionic strength has been fixed at 1.0 mM. Surface equilibrium constants and surface site densities listed in Table I1 are used in the calculation; results are shown in Table VII. As the oxide surface becomes more basic (A1203> FeOOH > TiOz), the surface charge density arising from protonation becomes substantially higher, and long-range electrostatic forces favoring nonspecific adsorption of anions increase in magnitude. Thus, nonspecific adsorption should be considered for all A1203 and FeOOH suspensions under low ionic strength conditions (below M). The development of surface charge on TiOz, in contrast, is substantially lower. Under most conditions, nonspecific adsorption onto Ti02 can probably be ignored. Other Pertinent Issues. Up to this point, we have discussed how reaction stoichiometry can be discerned from experimental observations. We have also discussed ways of estimating the magnitude of long-range electrostatic forces arising from the electrical double layer. There are additional phenomena influencing the extent of adsorption, however, which are more difficult to assess. Consider the first of the assumptions made in our modeling effort: all surface sites are chemically identical. In the case of crystalline oxides, crystallographic information can be used to estimate nearest-neighbor distances and coordinative arrangements on each cleavage plane. This approach reveals the presence of several structurally distinct surface groups, whose existence has been borne out using IR and other spectroscopic techniques (e.g., refs 30-32). Hiemstra et al. (28,29) have used this approach to identify surface-bound oxygen atoms coordinated to
one, two, and three metal atoms with oxide lattices and have estimated protonation constants for each type of site. It is likely that structurally distinct surface groups are also chemically distinct from one another. Despite this fact, uniform-site models are often able to predict changes in the extent of solute adsorption onto oxide surfaces known to contain several kinds of sites. It is important to point out that each subset of sites will obey a mass action expression. As long as the adsorbate concentration range being examined is not too broad, and as long as surface coverages are not too high, the aggregate behavior of the entire collection of sites will yield increases in adsorption that are proportional to increases in adsorbate concentration (Werner Stumm, personal communication). The second assumption used in our modeling effort is that all species adsorb at the same plane. Accordingly, all adsorbing ions of equal charge experience the same Poisson-Boltzmann electrostatic correction term. The triple-layer model, in contrast, assigns outer-sphere complexes to an adsorption plane which is some distance away from the adsorption plane for inner-sphere complexes and assigns a somewhat smaller electrostatic correction term. For many reasons, it is important to be able to distinguish inner-sphere from outer-sphere adsorption. Various spectroscopic methods offer the greatest potential in this regard; recent EXAFS work examining the adsorption of inorganic anions has been particularly informative (e.g., ref 33). Although stoichiometry alone cannot distinguish inner-sphere from outer-sphere adsorption, comparing the adsorption behavior of structurally related adsorbates can be quite useful. Geometric effects, including bond lengths and bond angles, should have a substantially greater effect on inner-sphere complex formation, especially when chelate formation can take place. Outer-sphere complexes, in contrast, should be relatively insensitive to small changes in ligand structure; only changes in the charge and radius of the ligand donor molecule should have a significant affect on outer-sphere adsorption. Specificadsorption of electrolyte ions represents a major source of uncertainty in our modeling efforts. In our attempt at evaluating the magnitude of nonspecific adsorption, it was assumed that surface charge was directly related to surface protonation level. If, however,significant adsorption of C1- onto positive FeOOH surfaces takes place, the lowering of surface charge (and surface potential) would lessen the accumulation of counterions (e.g., ligands) and lessen the repulsion of co-ions (e.g., protons). Under these conditions, the calculations used to generate Figure 2 would yield overestimates. Finally, assumptions have been made that the activity coefficients of dissolved species near the oxide/water interface are identical to those of species in bulk solution and that adsorbed species have activity coefficients equal to 1.0. These assumptions are provisional ones, since considerable evidence exists that important solvent properties, including dielectric constant, are perturbed near charged interfaces (34, 35). These changes in solvent properties near surfaces should alter the values of activity coefficient corrections. Additional information is needed before these and other secondary effects involving the oxidelwater interface can be accounted for. Previous work examining the adsorption of carboxylic acids, phenols, anilines, and pyridines at the oxidelwater interface (4,12-14) largely ignored the nonspecific contribution toward adsorption. This approach is justified Environ. Scl. Technoi., Vol. 27, No. 5, 1003
001
when experiments are restricted to high ionic strength conditions (e.g., 0.10 M NaC104) or when oxides bearing low surface charge density (e.g., Si021are employed. Two and sometimes three adsorption stoichiometries, each with a corresponding surface complex formation constant, have often been used to provide agreement with experimental data. Kummert and Stumm (12)employed a single plane of adsorption as discussed here, while Zachara et al. ( 4 ) employed two separate planes of adsorption as part of the triple-layer model.
Materials and Methods All solutions and suspensions were prepared from 18 MQ cm resistivity water (DDW, Millipore Corp.). All and rinsed several glassware was cleaned in 5 N "03 times with DDW prior to use. Inorganic reagents were of analytical grade, and the organic reagents were used as received. Stock organic solutions weremade slightly acidic through the addition of HCl to minimize autoxidation and used within 24 h of preparation. Acidity constants were determined from changes in UV absorbance during acid/ base titrations. 2-AP and 4-methyl-2-AP slowly degraded in the presence of 0 2 . Sparging stock solutions and oxide suspensions with argon was found to minimize oxidative degradation. Particle Preparation and Characterization. y-Muminum oxide (A1203, type C), amorphous silica (Si02, Aerosil200),and titanium dioxide (TiOz,primarily anatase, type P25) were obtained from Degussa Corp. and used without purification. FeOOH was synthesized by following a modification of the method of Atkinson et al. (36) in which 0.25 M Fe(N03)2 was slowly pumped into a 0.50 M KOH stock solution at 5 "C and then placed in a 70 "C constant-temperature oven for 24 h. FeOOH was repeatedly rinsed by centrifugation, removal of supernatant, and resuspension and then was freeze-dried. Torrents (37) provides details concerning characterization of the oxide samples. The specific surface area (S, m2/g)was determined according to the BET method (38). The number of fluoride-binding surface sites (mol/m2) was determined following the method of Sigg and Stumm (27). One gram of each oxide was added to 100 mL of 5.0 mM NaF solution, and the pH was adjusted to 4.7. After an adsorption period of 3 h, the remaining free fluoride in the suspension was determined using a fluoride-specific ion electrode (Orion Corp.). Monolayer coverage (one Fatom adsorbed on each exchangeable surface site) was assumed. Acid/base titrations were performed in a 15-mLdoublewalled beaker which was connected to a 25 OC water bath and sparged with N2 to eliminate C02. Suspensions were titrated by adding 5-100 p L of 0.1 N NaOH at intervals of 3 min. The pH,,, (pH of zero proton condition) of each sample was determined by conducting titrations at different electrolyte (NaN03)concentrations. Characteristics of these oxides are summarized in Table 11. Acidity constants and aK4intr) were obtained from graphical extrapolation of transformed titration data to zero surface charge conditions followingthe procedure outlined for the diffuse-layer model by Huang and Stumm (39)and Huang (40). Specific adsorption of electrolyte ions (reactions 5 and 6 in Table 111)was ignored in the calculation of surface acidity constants. Adsorption Experiments. Predetermined amounts of oxide and DDW were added to 30-mL amber-tinted 902
Environ. Sci. Technol., Vol. 27, No. 5, 1993
glass vials sealed with Teflon/silicon rubber septa. Suspensions were continuously stirred in a 25 "C constant temperature bath using Teflon-coated magnetic stir bars. Suspension pH was adjusted by adding small amounts of HCl or NaOH, and the ionic strength was set through by adding NaCl or NaN03. In some cases, 1.0-2.5 mM concentrations of buffer (acetate or ammonia) were added to the suspensions. For all the organic compounds examined, the extent of adsorption measured after 10min equilibration with oxide suspensions equaled the extent of adsorption measured 1 h or more later. Once this was determined, samples were allowed 15-30 min of contact before samples were removed and analyzed. This contact period may not be sufficient to yield true "equilibrium" adsorption values, since diffusion into pores within the oxide may require time scales of several days or more. The experiments do, however, represent a self-consistent set of adsorption data which can be used to test our understanding of the processes involved. TiOz, A&, and Si02 were removed from suspension using an Eppendorf centrifuge (15 000 rpm for 7 min), while FeOOH particles were removed by filtration (0.2 pm Nuclepore Corp. filter). 2-PM concentrations in supernatant solution were measured using UV spectrophotometry. Proportional decreases in absorbance were observed at three or more separate wavelengths, providing evidence that adsorptive loss, not chemical alteration, was responsible for the spectral changes. Concentrations of all other test compounds were measured in supernatant solution using reversed-phase HPLC (p-Bondapak-Cls column, Waters Corp.) with UV detection. Blank experiments were performed to ensure that test organic compounds did not adsorb to the polypropylene and polyethylene centrifuge tubes or to the polycarbonate Nuclepore filters. The difference between test compound concentrations in supernatants collected from oxide suspensions and in particle-free solutions was used to calculate the extent of adsorption. In order for this method to be valid, it must be shown that significant test compound degradation does not take place. Supernatant concentrations of test compounds were monitored as a function of time. A rapid decrease in supernatant concentration followed by imperceptible changes for extended periods of time provides indirect evidence that degradation is not taking place. Three methods for the direct recovery of test compounds from oxide surfaces were also considered: (i) addition of miscible organic solvents that lessen hydrophobic contributions to adsorption, (ii) addition of a second adsorbable compound that can displace the test compound, and (iii) pH adjustment to a range where adsorption is low. The first method was found insufficient for many of the test compounds examined (log KO,values for the neutral species are low). The second method is unreliable because competitive adsorption processes are not understood with sufficient certainty. We found that the third method worked reliably in many instances. Complete recovery of 2-aminophenols,for example, could be achieved by sharply lowering the pH to 2.
Results and Discussion Adsorption experiments involving the ligands listed in Table I and the oxides listed in Table I1 will now be
described. 2-Aminophenols and 2-pyridinemethanol are considered bidentate ligands, Although nitro substituents can participate in complex formation, their contribution is relatively weak, Thus, 2,4-dinitrophenol(2,4-DNP) and 2,6-dinitrophenol (2,6-DNP) are expected to behave as monodentate ligands. Our objective is to evaluate whether a single choice of proton stoichiometry selected from reactions 7-10 in Table I1 can adequately account for the effect of pH and ionic strength on organic ligand adsorption. Surface acidity constants used in the diffuse-layer model were determined separately using acid/base titrations, and no attempt was made to account for the adsorption of the supporting electrolyte. Thus, the intrinsic surface equilibrium constant for organic ligand adsorption remains the single fitting parameter in the modeling exercise. The intrinsic surface equilibrium constant was selected by systematically changing its value until the amount adsorbed as a function of pH matched most closely the experimental data. Adsorption of 2,4- and 2,B-Dinitrophenol. Preliminary experiments in our laboratory indicated that the adsorption of phenol (pK, 9.98) and more basic phenols such as 2,bdimethylphenol (pK, 10.40), 2,6-dimethylpheno1 (PKa 10.62),3,5-dimethylphenol (PKa 10.20),and 2,4,6trimethylphenol (pKa10.99)onto TiOz, A1203, and FeOOH surfaces was below the detection limit of our analytical methods. All these compounds exist as the protonated, neutral species HLo at pH values below the pH,,, of the oxides used in the experiments. 2,4-Dinitrophenol(2,4-DNP) is substantially less basic than the phenols listed above; it possesses two electronwithdrawing nitro substituents which lower the pKa to 4.11, below the pH,,, of the test oxides. We will first examine the adsorption of 2,4-DNP (5.88 X M)in 10 g/L Ti02 suspensions. Because the oxide loading is high (STequal to 2.74 X M), 2,4-DNP is capable of occupying only a small fraction of sites and, therefore, has a negligible effect on surface charge. The extent of 2,4DNP adsorption is shown in Figure 3 under low ionic strength conditions (ionic strength near 1.0 mM, arising from use of a 3.0 mM acetate buffer). Surface charge and 2,4-DNP protonation level are also shown in Figure 3, calculated from property data listed in Tables I and I1 and the diffuse-layer model. 2,4-DNP adsorption is negligible at neutral and alkaline pH but rises to substantial levels near and below the pKa. The diffuse-layer model can be employed to simulate the adsorption data. We first examine the adsorption stoichiometry expressed by reactions 7 and 8 in Table 111. Since the model treats these two equations in a mathematically equivalent fashion, >MOH2+L- and >ML are indistinguishable from one another. Using the ligand mass balance equation presented previously in Table VI, Part C, a value of log ,K7,aintrcan be selected that matches most closely the experimental data. As Figure 3 demonstrates, the stoichiometry of reactions 7 and 8 very accurately predicts the sharp increase in adsorption occurring as the pH is decreased from 5.5 to 4.5. Alternative protonation stoichiometries (i.e., reactions 9 and 10)yield model results that do not match the experimental data. Because the accumulation of positive charge on the Ti02 surface is low within the pH range studied, nonspecificadsorption arising from long-range electrostatic interactions is insignificant. Using the results from Table V, the stoichiometry
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represented by reactions 7 and 8 can be examined in greater detail: The pH behavior of the three terms comprising this equation can be summarized: (i) [H+]bulkincreases by a factor of 10 for every unit decrease in pH; (ii) when pH > pK,, [L-lbulk maintains a constant, high value; (iv) [>MOHl maintains a relatively constant value. The product of these terms, which determines the value of [L]&, growsto its maximum value as the pH is decreased from (PKa + 1.0) to pKa. The pK, values for the two isomers 2,4-DNP and 2,6DNP are quite similar. Adsorption of 2,6-DNP (not shown) as a function of pH is virtually identical to the adsorption of 2,4-DNP, including the sharp increase in adsorption as the pH is decreased from 5.5 to 4.5. These observations are a strong indication that the nature of the adsorption reaction is very similar for the two isomers. Experimental measurements presented in Figure 4 reveal a slight to moderate decrease in the extent of 2,4DNP adsorption as the electrolyte concentration is increased to 0.10 M through the addition of NaC1. The diffuse-layer model can be applied to the high electrolyte concentration case, using the same value of log aK7,8intr Environ. Scl. Technol., Vol. 27, No. 5, 1093 003
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selected earlier to match the low electrolyte concentration case. It is important to point out that this model does not consider the specific adsorption of electrolyte ions. As the figure indicates, the diffuse-layer model actually predicts a slight increase in the maximum extent of adsorption as the ionic strength is increased. Unless specificadsorption of electrolyte ions takes place, the effect of increasingthe electrolyte concentration should be independent of the electrolyte employed. Figure 5 indicates, however, that the addition of NaN03 does not diminish the extent of 2,4-DNP adsorption as much as the addition of NaC1. At a concentration of 0.10 M, for example, there is a 22.3% difference in the extent of 2,4DNP adsorption in the presence of the two electrolytes. Thus, specific interaction of C1- with Ti02 surface sites is taking place. Specific adsorption of NOa-cannot be ruled out; we only know that it is less than the specific adsorption of c1-. Specific adsorption of electrolyte ions can be accounted for by including values for aKgintr and *Kgintrfrom Table I11 in the diffuse-layer model. Although reliable values for these constants are not available at present, the effect on model results is clear: as the electrolyte concentration is increased, competitive adsorption of electrolyte ions lowers the extent of ligand adsorbed. Next, we consider adsorption of 2,4-DNP onto FeOOH. Unlike TiO2, a moderate positive charge develops on the SO4
Environ. Sci. Technol., Vol. 27, No. 5, 1993
FeOOH surface below pH 7.0. As Figure 6A indicates, there is a distinct adsorption maximum near pH 4.5, and adsorption diminishes substantially when the electrolyte (NaC1) concentration is increased. A diffuse-layer model following the procedure outlined in Table VI, Part C, was used to sum contributions from both nonspecific and specific adsorption. A value of log aK7,aintr equal to 6.45 was selected to yield an adsorption maximum equal to 38 % ,reached by a single point in Figure 6A. Given the considerable variation observed in the data, however, a lower value of log aK7,8intr would perhaps be more appropriate. Results of the diffuse-layer model are shown in Figure 6B,C. At low electrolyte concentrations, 2,4-DNP adsorption onto FeOOH arises from both nonspecific interactions and from surface complex formation. Considering the variability in the experimental data, the nonspecific contribution to adsorption under these conditions lies somewhere between 20 and 30%. At high electrolyte concentration, the contribution arising from nonspecific adsorption sinks to negligible levels. Finally, we consider the adsorption of 2,4-DNP onto A&. As we have discussed earlier, A1203exhibits a pH,,, that is substantially higher than the other two oxides and acquires a higher surface charge density under acidic conditions. Figure 7A indicates that the adsorption of 2,4-DNP onto 10 g/L A1203reaches a sharp maximum near pH 5.5 at low ionic strength and decreases substantially as the NaCl concentration is increased to 0.10 M. We attempted to model the data by varying the value of log aK7,aintruntil the adsorption maximum predicted by the model matched the maximum observed in the experimental data. As shown in Figure 7B, the model slightly overestimated the extent of adsorption at low electrolyte concentration, even when surface complex formation was omitted from the model. Thus, nonspecific adsorption arising from long-range electrostatic forces can entirely explain the 2,4-DNP adsorption data. The model also predicts that adsorption drops to negligible levels at high electrolyte concentration. Comparisonsconcerningthe adsorption of 2,4-DNP onto the three oxides leads to the following conclusions. Because the chargingof the Ti02 surface is low, nonspecific adsorption is quite negligible, and surface complex formation is the dominant adsorption mechanism. In the case of FeOOH,moderate development of a positive surface charge takes place within neutral to acidic pH conditions; both nonspecific adsorption and surface complex formation are important processes. A1203 acquires the highest positive surface charge of the three oxides considered. Nonspecific adsorption is so strong relative to surface complex formation that it entirely dominates the adsorption of 2,4-DNP onto Al2O3. In the case of Ti02 and FeOOH, specific adsorption consumes one proton for every 2,4-dinitrophenolate anion adsorbed, consistent with the formation of either the innersphere complex >ML or the outer-sphere complex >MOHz+L-. The placement of two bulky nitro substituents ortho to the donor group should impede entry of 2,6-DNP into the inner-coordination sphere of surfacebound metal atoms. Close similarities observed in the adsorption of 2,4-DNP and 2,6-DNP onto Ti02 provide indirect evidence that outer-sphere rather than innersphere complex formation is taking place. Indirect evidence also indicates that specific adsorption of C1- onto the Ti02 surface is taking place. This
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phenomenon leads to a greater decrease in 2,4-DNP adsorption than the diffuse-phase model predicts as the NaCl concentration is increased. This finding indicates that our estimates of nonspecificadsorption in the presence of NaCl are too high; specific adsorption of C1- will lower the positive surface charge, diminishing the magnitude of long-range electrostatic forces. Adsorption of Substituted 2-Aminophenols. Preliminary experiments in our laboratory indicated that 2-aminophenols adsorb at significant levels onto Ti02 surfaces, while 3-aminophenols do not. The obvious difference between these two sets of isomers is that the phenolate and aromatic amine substituents of 2-aminophenols can form chelates with dissolved and surfacebound metal atoms, while 3-aminophenols cannot. In our discussions of the adsorption of 2-aminophenols, the consequences of surface chelate formation must be dealt with explicitly. Most importantly, surface chelate formation affects the protonation stoichiometry of the adsorption process. The protonation stoichiometry, in turn, sets the pH dependence of the adsorption process. FeOOH was not included in our adsorption experiments because oxidation of adsorbed 2-aminophenols by surfacebound Fe(II1) was expected to take place. All of the 2-aminophenols examined adsorbed to a significant extent onto 10 g/L TiO2, while adsorption onto 10 g/L A1203 was below the sensitivity limit of our analytical method. Adsorption onto A1203 is not limited by the availability of surface sites. According to our F- adsorption measurements, ST in 10 g/L A1203 suspensions is 1.7 times higher than in 10 g/L Ti02 suspensions.
The lack of observable adsorption onto A1203 is most readily explained by the selectivity of ligand donor groups toward different metals. Oxygen-donor atoms are classic "hard" ligands which form strong complexes with metals possessing high charge-to-radius ratios such as Al(II1) (0.53-A atomic radius) and Ti(1V) (0.60-A atomic radius) (41). Nitrogen-donor atoms, while still classified as "hard" ligands (42),are more polarizable than oxygen-donoratoms and, therefore, show a preference for transition metal ions (43). Thus, greater stability of the amine-Ti(1V) bond relative to the amine-Al(II1) bond raises the extent of 2-aminophenol adsorption onto Ti02 to within measureable levels. We will now examine the adsorption of 2-aminophenols onto 10 g/L Ti02 in greater detail. The surface charge arising from surface protonation and the protonation equilibria of 2-AP in bulk solution at high electrolyte concentration (0.10 M NaCI) are shown in Figure 8A,B. Experimental measurements for the adsorption of 4.84 X M 2-AP are presented in Figure 8C, along with results from the diffuse-layer model. The model was performed using the protonation stoichiometry expressed by reaction 8 in Table I11 (one proton consumed per mole of aminophenolate anion adsorbed, yielding a neutral surface complex). was adjusted until the adsorption maximum predicted by the model equaled the adsorption maximum determined by experiment. The model results and the experimental data are in close agreement. As the pH is increased, the extent of adsorption rises to a plateau above pH 4.7, which corresponds to the pK, for the aromatic amine substituent. In agreement with our study Envlron. Sci. Technol., Vol. 27, No. 5, 1993 906
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of 2,4-DNP, adsorption reaches a maximum value within a pH range where the neutral form of the test organic compound (HLO) dominates bulk solution speciation. Despite argon sparging, small amounts of 02 entering the suspension interfered with measurements of 2-AP adsorption above pH 8. For this reason, the model prediction that adsorption should decrease substantially above pH 9 could not be tested. It is important that alternative protonation stoichiometries for the adsorption process be considered. Figure 9 shows the high electrolyte concentration adsorption data along with diffuse-layer model calculations using adsorption stoichiometries represented by reactions 8-10 from Table 111. Reaction 9 involves consumption of two protons for each aminophenolate anion adsorbed, yielding a surface complex bearing a +1 charge. Reaction 10 does not bring about proton consumption; it yields a surface complex bearing a -1 charge. Reaction 8 clearly provides the best representation of the adsorption data, since model results from reactions 9 and 10 predict an adsorption maximum arising at pH values too acidic or too basic. Models based upon reactions 9 and 10 exhibit a substantial decrease in adsorption as the electrolyte concentration is decreased. 906
Environ. Sci. Technol., Vol. 27, No. 5, 1993
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Flgure 10. Adsorption of 2-aminophenols onto 10 g/L TiOn: (A) 2-AP, 0.10 M NaCI; (B) 2-AP, low ionic strength (near 1.0 mM); (C) 4-nitro2-AP, 0.05 M NaCI; (D) 4-nitro-2-AP, low ionic strength. Filled symbols represent the experimental adsorption data, while continuous curves represent diffuse-layer model results. Test compound concentraM 2-AP; 4.46 X M 4-nitro-2-AP. tions: 4.84 X
This always occurs when the postulated reaction converts surface sites bearing no net charge into positively or negatively charged surface complexes. 2-AP adsorption onto Ti02 in suspensions containing high (0.10 M NaC1) and low (ionic strength near 1.0 mM) electrolyte concentrations are presented in Figure 10A,B. Experimental results show a slight increase in 2-AP adsorption when the electrolyte concentration is decreased; no shift in the pH dependence is observed. When the surface complex formation constant (*Ksintr)determined under high ionic strength is employed, the model predicts a slight decrease in 2-AP adsorption under low electrolyte concentration conditions. Competition between C1- and 2-AP for Ti02 surface sites can explain this discrepancy. If the extent of C1-
adsorption was known, the lowering of the concentration of free surface sites could be considered in the calculation Of increasing its value. The higher value of a&intr, in turn, would yield a higher extent of adsorption from the model under low electrolyte concentration conditions and greater agreement with the experimental data. Experimental and model4-nitro-2-AP adsorption results from suspensions containing high (0.05 M NaC1) and low (ionic strength near 1.0 mM) electrolyte concentrations are presented in Figure 10C,D. The same stoichiometry for surface complex formation (reaction 8 in Table 111) was used to model adsorption. 4-Nitro-2-AP is not very susceptible toward autoxidation, so reliable adsorption measurements can be made under alkaline, neutral, and acidic conditions. Both the experimental data and the diffuse-layer model exhibit a sharp increase in adsorption as the pH is increased in the vicinity of pK,1 (at pH 3.1), a plateau region near neutral pH where HLO is the predominant species in bulk solution, and a sharp decrease in adsorption in the vicinity of pKa2 (at pH 7.6). This sharp decrease at high pH displayed by the experimental data is important, since it confirms the validity of the adsorption stoichiometry selected in this range, which we were unable to do with 2-AP. Comparing the adsorption of 2-AP and 4-nitro-ZAP provides our first opportunity to explore substituent effects. The electron-withdrawing nitro substituent substantially lowers the basicity of the aromatic amine and phenolate donor groups, reflected in a lowering of pKa1 and pKa2. As shown in Figure 10, there is a corresponding shift in the plateau of maximum adsorption toward more acidic pH, exhibited by both the experimental data and the diffuse-layer model results. Values of ,Ksintr used to model the adsorption data involving 2-aminophenols are presented in Table I. The value of the surface complex formation constant for 2-AP is times higher than the corresponding value for 4-nitro-2-AP. Thus, log K for surface complex formation increases as the ligand basicity increases. This is the expected result, since substituents that increase the ligand donor ability toward protons (i.e., basicity) should also increase the ligand donor ability toward metal ions. As demonstrated in Table I, the other 2-aminophenols included in this study (4-methyl-2-AP and 3-amino-2NAPH) also obey this trend. For this class of compounds, a plot of log versus pKa2 is linear. Despite the fact that for 4-nitro-2-AP is 101.9times lower than that for 2-AP, the maximum extent of adsorption is 2.8 times higher. This is a direct consequence of the mass balance equation for ligand (TOTL) presented in Table V. Because surface-bound metal atoms must compete with protons for binding the ligand, it is the magnitude Of relative to aK1and “K2(proton binding constants defined in Table 111)which is most important in determining surface coverage. Other secondary substituent effects may influence the extent of adsorption. It was mentioned earlier that the plot of log versus pKa2 is linear; 2-aminophenols bearing hydrophobic substituents (e.g., 3-amino-2-NAP) do not behave significantly different from those that do not. It would be fruitful to examine whether the linear relationship between log aK8intr and pKancontinues to apply for 2-aminophenols bearing even more hydrophobic substituents, such as 1-amino-2-hydroxyanthracene.
The experiments effectively demonstrate that chelate formation between the ligand donor groups and surfacebound Ti(1V)atoms is the dominant process bringing about the adsorption of 2-aminophenols. The stoichiometry of the adsorption reaction is also straightforward one proton is consumed for every molecule of aminophenolate anion adsorbed. Adsorption of 2-Pyridinemethanol. 2-PM contains a nitrogen-donor group and an oxygen-donor group capable of chelating dissolved and surface-bound metals. As with 2-AP, the nitrogen-donor group imparts selectivity to the adsorption process; appreciable 2-PM adsorption onto Ti02 takes place, while adsorption onto A1203 is below the sensitivity limit of our analytical method. Both 2-PM and 2-AP form five-membered chelate rings, and the basicity of the nitrogen-donor groups is virtually the same. The sp3 nature of the nitrogen atom in 2-AP and the sp2 nature of the nitrogen atom in 2-PM is expected to affect chelate formation, however. The pyridyl nitrogen of 2-PM has lost electron density via resonance with the aromatic ring, and its entry into the inner-coordination sphere of metals experiences greater steric hindrance from ortho substituents. Perhaps the greatest difference between these two adsorbable compounds is the nature of the oxygen-donor groups. As evident from the pK, values listed in Table I, the alcoholic OH group of 2-PM is substantially more basic than the phenolic OH group of 2-AP. Figure 11A shows the adsorption of 5.08 X M 2-PM onto 7.0 g/L Ti02 in high electrolyte concentration (0.11 M NaC1) medium. The filled symbols represent the experimental data, while the continuous curves are diffuselayer model simulations based upon adsorption stoichiometries expressed in reactions 7-9 of Table 111. For each stoichiometry, an equilibrium constant was selected that yielded an adsorption maximum of comparable magnitude to the experimental data. Reaction 8 stoichiometry, yielding a neutral surface complex, predicts the rise in adsorption as the pH is increased in the vicinity of pKal reasonably well but overestimates the amount adsorbed at high pH. Reaction 9 stoichiometry, yielding a +LOcharged surface complex, also predicts the rise in adsorption near pKal reasonably well. Reaction 9 does, however, drastically underpredict the amount adsorbed as the pH is increased. Reaction 10 stoichiometry, yielding a -1.0charged surface complex, yields the worst fit; it predicts that the highest amounts of adsorption take place under alkaline conditions and that the extent of adsorption near neutral pH is nearly negligible. Surface equilibrium constants selected to match the high electrolyte concentration experiments have been used to estimate adsorption behavior at low electrolyte concentrations (Figure 11B). The experimental data yield a very slight dependence on pH; the adsorption maximum shifts to higher pH, with very little change in height. Reaction 8 stoichiometry is very similar in this regard; it predicts a slight-to-moderate decrease in adsorption as the electrolyte concentration is decreased and a slight widening of the plateau of maximum adsorption. At this point, reaction 9 and 10 stoichiometries can be discarded, since they both predict a much higher electrolyte concentration effect than seen in the experimental results. Although reaction 8 stoichiometry provides the best fit to the experimental observations concerning 2-PM adsorption, the fit is not nearly as good as those accomplished Environ. Scl. Technol., VoI. 27,
No. 5, 1993 007
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with 2,4-DNP and 2-AP. Viewing the fitting procedure mathematically, the experimental observation seem to match an organic compound with the same pKal value as 2-PM but with a substantially lower value of pKa2. At this point, it is tempting to evoke the complexity of the adsorption process to explain the discrepancy between the 2-PM experimental observations and the model fit. Uncertainties arising from (i) structure-chemical differences among surface sites, (ii) geometrical and steric constraints affecting surface complex formation, and (iii) poorly understood charging behavior of oxide surfaces at high pH are difficult to assess at this point. I t should be kept in mind, however, that these models are all based upon a single adsorption stoichiometry and a single fitted parameter, the surface complex formation constant. Despite the simplicity of the modeling effort, the agreement with experimental observations is impressive. Conclusions The adsorption of organic ligands at the oxidelwater interface arises from a combination of long-rangeprimarily electrostatic interactions and near-range physical and chemical interactions. The program HYDRAQL (24), based upon the complex equilibrium program MINEQL (23) accounts for these interactions using a diffuse-layer model for the electrical double layer and a surface comprised of chemically identical sites. The effects of pH and ionic strength on ligand adsorption are a logical consequence of (i) stoichiometries of surface complex formation reactions and the magnitude of corresponding surface complex formation constants, (ii) mass balance equations for ligand and for surface sites, and (iii) PoissonBoltzmann terms which relate the concentrations of charged species at the oxidelwater interface to concentrations in bulk solution. 2,4-DNP forms surface complexeswith Ti02 and FeOOH sites which can be modeled with a single choice of reaction stoichiometry: one proton is consumed for every phenolate anion adsorbed, leaving the oxide surface charge unaltered. Long-range electrostatic interactions, which play a negligible role in 2,4-DNP adsorption onto Ti02, increase in importance in going from a moderately charged oxide surface such as FeOOH to a highly charged surface such as A1203. 2-Aminophenols adsorb onto Ti02 surfaces via inner-sphere surface chelates and obey the following stoichiometry: one proton is consumed for every ami908
Environ. Sci. Technol., Vol. 27, No. 5, 19g3
nophenolate anion adsorbed. Comparisons among four substituted 2-aminophenols reveal a linear relationship between surface complex formation constants and ligand basicity. In the case of 2-PM, our ability to model experimental data using a single adsorption stoichiometry was not as successful as with the other ligands examined. Adsorption of organic ligands at oxidelwater interfaces is an inherently complex process. Despite this complexity, models based upon a single adsorption stoichiometry and a single fitting parameter, the intrinsic surface equilibrium constant, are surprisingly successful at predicting the effects of changing pH and electrolyte concentration on adsorption. Experiments of this kind, in combination with new spectroscopic techniques, should bring us closer to our ultimate goal: the ability to predict the extent of organic ligand adsorption based upon the structure and properties of the organic ligand, upon the structure and properties of the oxide surface, and upon the composition of the supporting aqueous medium. Acknowledgments This research was supported by the U.S. Geological Survey (USGS), Department of the Interior, under USGS Award 14-08-0001-G1647. The views and conclusions contained in this document are those of the authors and should not be interpreted as necessarily reflecting the official policies, either expressed or implied, of the US. Government. Eric Weber, George Bailey, and other members of the Chemistry Branch, USEPA Environmental Research Laboratory in Athens, GA, provided advice and technical assistance, which are gratefully acknowledged. Prof. Charles O'Melia (Johns Hopkins University) also provided useful input. Literature Cited (1) Karickhoff, S. W.; Brown, D. S.; Scott, T. A. Water Res. 1979,13,241-248. (2) Chiou, C. T.; Peters, L. J.; Freed, V. H. Science 1979,206, 831-832. (3) Brownawell, B. J.; Chen, H.; Collier, J. M.; Westall, J. C. Enuiron. Sci. Technol. 1990, 24, 1234-1241. (4) Zachara, J. M.; Ainsworth, C. C.; Cowan, C. E.; Schmidt, R. L. Enuiron. Sci. Technol. 1990, 24, 118-126. ( 5 ) Stone, A. T. Enuiron. Sci. Technol. 1987,21, 979-989. (6) LaKind, J. S.;Stone, A. T. Geochim.Cosmochim. Acta 1989, 53, 961-971.
(7) Laha, S.;Luthy, R. G. Environ. Sci. Technol. 1990,24,363373. (8) Waite, T. D.; Torikov, A. J . Colloid Interface Sci. 1987, 119, 228-235. (9) Waite, T. D.; Torikov, A.; Smith, J. D. J . Colloid Interface Sci. 1986,112,412-420. (10) Torrents, A,; Stone, A. T. Environ. Sci. Technol. 1991,25, 143-149. (11) Torrents, A.; Stone, A. T. Catalysis of picolinate ester hydrolysis a t the oxidelwater interface: inhibition by coadsorbed species. Submitted to Environ. Sci. Technol. (12) Kummert, R.; Stumm, W. J . Colloid Interface Sci. 1980, 12, 119-216. (13) Davis, J. A,; Leckie, J. 0. Environ. Sci. Technol. 1978, 12, 1309-1315. (14) Balistrieri, L. S.; Murray, J. W. Geochim. Cosmochim. Acta 1987,51, 1151-1160. (15) Westall, J. C.; Hohl, H. Ado. Colloid Interface Sci. 1980,12, 265-294. (16) Castellan, G. W. Physical Chemistry; Addison-Wesley: Reading, MA, 1971. (17) Bard, A. J.; Faulkner, L. R. Electrochemical Methods, 1st ed.; Wiley-Interscience: New York, 1980; pp 501-507. (18) van Olphen, H. Anlntroduction to Clay Colloid Chemistry, 2nd ed.; Wiley-Interscience: New York, 1977. (19) Stumm, W.; Morgan, J. J. Aquatic Chemistry; WileyInterscience: New York, 1981. (20) Martell, A. E.; Smith, R. M. Critical Stability Constants; Plenum: New York, 1977; Vol. 3. (21) Martell, A. E.; Smith, R. M. Critical Stability Constants; Plenum: New York, 1982; Vol. 5. (22) Smith, R. M.; Martell, A. E. Critical Stability Constants; Plenum: New York, 1989; Vol. 6. (23) Westall, J. C.; Zachary, J.;Morel, F. MINEQL: A computer program for the calculation of chemical equilibrium composition of aqueous systems. Technical Note No. 18, Ralph M. Parsons Laboratory, MIT, Cambridge, MA, 1976. (24) Papelis, C.; Hayes, K. F.; Leckie, J. 0. HYDRAQL: A computer program for the computation of chemical equilibrium composition of aqueous batch systems including surface-complexation modeling of ion adsorption at the oxide/solution interface. Technical Report No. 306, Department of Civil Engineering, Stanford University, Stanford, CA, 1988.
(25) Dzombak, D. A.; Morel, F. M. M. Surface Complexation Modeling: Hydrous Ferric Oxide;Wiley-Interscience: New York, 1990. (26) Westall, J. C. In Aquatic Surface Chemistry; Stumm, W., Ed.; Wiley-Interscience: New York, 1987; Chapter 1. (27) Sigg, L.; Stumm, W. Colloids Surf. 1981, 2, 101-117. (28) Hiemstra, T.; van Riemsdijk, W. H.; Bolt, G. H. J . Colloid Interface Sci. 1989, 133, 91-104. (29) Hiemstra, T.; DeWit, J. C. M.; van Riemsdijk, W. H. J. Colloid Interface Sci. 1989, 133, 105-117. (30) Parfitt, R. L.; Farmer, V. C.; Russell, J. D. J. Soil Sci. 1977, 28, 29. (31) Parfitt, R. L.; Fraser, A. R.; Russell, J. D.; Farmer, V. C. J. Soil Sci. 1977, 28, 40. (32) Parfitt, R. L.; Fraser, A. R.; Farmer, V. C. J . Soil Sci. 1977, 28, 297. (33) Hayes, K. F.; Roe, A. L.; Brown, G. E.; Hodgson, K. 0.; Leckie, J. 0.;Parks, G. A. Science 1987,238, 783-785. (34) Nurnberg, H. W. In Membrane Transport i n Plants; Zimmerman, U., Dainty, J., Eds.; Springer-Verlag: New York, 1974. (35) Nurnberg, H. W.; Wolff, G. J . Electroanal. Interfacial Electrochem. 1969, 21, 99-122. (36) Torrents, A. Ph.D. Thesis, Johns Hopkins University, Baltimore, MD, 1992. (37) Atkinson, R. J.; Posner, A. M.; Quirk, J. P. J. Phys. Chem. 1967, 71. 550. (38) Brunauer, S.;Emmett, P. H.; Teller, E. J. Am. Chem. SOC. 1938, 60, 309. (39) Huang, C. P. In Adsorption of Inorganics at the Solid/ Liquid Interface; Anderson, M. A., Rubin, A. J. Eds.; Ann Arbor Science: Ann Arbor, MI, 1981. (40) Huang, C. P.; Stumm, W. J . Colloid Interface Sci. 1973,43, 409-420. (41) Bell, C. F. Principles and Applications of Metal Chelation; Clarendon Press: Oxford, England, 1977; p 33. (42) Pearson, R. G. Science 1966, 151, 172-177. (43) Evers, A.; Hancock, R. D.; Martell, A. E.; Motekaitis, R. J. Inorg. Chem. 1989,28, 2189-2195.
Received for review July 21,1992. Revised manuscript received January 8, 1993. Accepted January 20, 1993.
Envlron. Sci. Technol., Vol. 27,
No. 5, 1993 909
