Portland and Eugene Airsheds”, final report to the Oregon Department of Environmental Quality, May 1979. (8) Clayton, G. C.; Arnold, J. F.; Patty, F. A. Science 1955, 122, 751. (9) Lodge, J. P.; Bien, G. S.; Suess, H. E. Int. J . Air Pollut. 1960,2, 309. (10) Currie, L. A. Proc. Conf. Radiocarbon Dating Accel. l s t , 1978 1978,372. (11) Currie. L. A.: Noakes. J. A,: Breiter. D. N. Proc. Int. Conf.Radiocarbon Dating, 9th 1979,158. (12) Currie, L. A,: MurDhv, - “ R. B. NES SDec. Publ. (U.S.)1977 No. 464,439. (13) Currie, L. A.; Kunen, S. M.; Voorhees, K. J.; Murphy, R. B.; Koch, W. F. “Analysis of Carbonaceous Particulates and Characterization of their Sources of Low-Level Radiocarbon Counting and Pyrolysis/Gas Chromatography/Mass Spectrometry”, Conference on CarbonaceousParticles in the Atmosphere, Conf-7803101 LBL, 1978. (14) Suess, H. E. Science 1955,122,415. (15) Lopes, J. S.; Pinto, R. E.; Almendra, M. E.; Machado, J. A. “Variation of 14CActivity in Portuguese Wines from 1940 to 1974”; Proceedings of the International Conference on Low-Radioactivitv Measurements and Applications,The High Tatras, Czechoslavakia, Oct. 1975,p 265.
(16) Olsson, I. U. Nobel Symp. 1970,12,17. (17) Stuiver, M.; Polach. H. A. Radiocarbon 1977,19,355. (18) Cooper, J. A.; Watson, J. G., Jr. J . Air Pollut. Control Assoc. 1980,30,10. (19) Hester, N. E. “Evaluation of Techniques to Determine the Impact of ParticulateMatter from Field and Slash Burning on Urban Areas”, draft final report to the Oregon Department of Environmental Quality,January 17,1979. (20) DeAneelis. D. G.: Ruffin. D. S.: Reznik. R. B. “Source Assessment: WoodlFired ’ Residential Combustion Equipment Field Tests”, U.S. Environmental Protection Agency report no. MRCDA-EPA-600/2-79-019,1979, draft copy. (21) Cooper, J. A. J. Air Pollut. Control. Assoc. 1980.30, 8. (22) Gove, H. Proc. Radiocarbon Dating Accel., 1st 2978,1978. (23) Muller, R. A. Science 1977,196,489.
Received for review August 1, 1979. Revised manuscript received February 23, 1981. Accepted May 7, 1981. Partial support by the Office of Environmental Measurements ( N B S ) ,the EPA EnergyEnvironmental Program (EPA-IAG-D6-E684),and the Oregon Department of Environmental Quality is gratefully acknowledged.
Conceptual Model for Metal-Ligand-Surface Interactions during Adsorption Mark M. Benjamin* Department of Civil Engineering, University of Washington, Seattle, Washington 98195
James 0. Leckie Department of Civil Engineering, Stanford University, Stanford, California 94305
Adsorption of trace metals from aqueous solution is dependent on pH, adsorbent and adsorbate concentration, and speciation of the metal in solution. In particular, complexation of metal ions by inorganic and organic ligands can dramatically increase or decrease adsorption compared to a ligand-free system. In this paper a semiquantitative model is presented to account for these changes. The model considers complexed species to be either “metallike” or “ligandlike”, depending on whether adsorption of the complex increases or decreases with increasing pH, respectively. Expected patterns for partitioning of the metal between the surface and solution are presented and shown to be qualitatively different for metallike and ligandlike complexes, particularly at low pH. This model can be combined with models for the charge-potential relationship in the electrical double layer to yield quantitative predictions. ~~~~~
~
~
Introduction
Whereas total analytical concentration of a chemical in an aquatic system was once thought to be an adequate measure of its significance, it has become increasingly clear that the impact of a given chemical constituent depends on its speciation. This is particularly true for heavy metals, which may be present in dissolved, colloidal, or particulate fractions, may form a pure solid phase or adsorb on the surface of a foreign particle, may associate with organic or inorganic material, forming covalently bonded molecules or “labile” or “nonlabile” complexes, and in many cases may be present in more 1050
Environmental Science & Technology
than one oxidation state ( I ) . Because the bioavailability and toxicity of heavy metals varies widely depending on their physicochemical form, it is important to understand how interconversions are controlled. One of the most fundamental of these interactions is partitioning between dissolved and surface (adsorbed) phases. There have been numerous studies of metal adsorption in “simple” systems, where the only reactions in which the metal ions participate are adsorption and hydrolysis (e.g., 2-5). These have consistently shown that pH is the dominant solution parameter controlling adsorption and that cation adsorption increases dramatically as solution pH increases. However, when complexing ligands are added to a system, the results cannot be generalized easily. Metal adsorption sometimes increases (6-8) and sometimes decreases (9-11) depending on the particular metal, ligand, adsorbent, and pH range being studied. In this paper, a conceptual and mathematical model is presented that can account a t least qualitatively for many of the observed effects of complexation on metal ion adsoiption. The model can be used in a predictive fashion by quantifying a term which describes the chargepotential relationship in the interfacial region. Several sets of equations have been proposed to describe this relationship (12-15). The model presented here is general and can be combined with any of the proposed interfacial charge-potential relationships to yield quantitative predictions. In a subsequent paper data for Cd adsorption onto four different solids in the presence of several complexing ligands are presented and analyzed in terms of the model ( 1 6 ) .
0013-936X/81/0915-1050$01.25/0
@ 1981 American Chemical Society
~~
Conceptual Basis of t h e Model
Systems containing metal, ligand, and solid can be viewed as consisting of two phases with metal-ligand equilibria reactions governing speciation in each, as pictured below:
Me”’
Table 1. Range of Adsorptive Characteristics Possible for Adsorption of Metal-Ligand Complexes pH-adsorption pattern of com pIex
similar to Me (“metallike”) similar to L (“ligandlike”) a
strength of MeL-surface bond a
(A) stronger than Me/S,IB)equal to Me/$, (C) weaker than Me/S (D) stronger than L/$,IE) equal to LIS,
(F)weaker than LIS
5 represents a surface-binding site.
MeLm-n
Ln-
8
adsor:eed
Table II. Equations Defining Adsorption Equilibria for Metallike and Ligandlike Complex Systemsa
solution phase
The bottom schematic of the surface represents a vacant site or, equivalently, a site occupied by a solvent molecule. The interaction between the metal and the ligand a t the surface may be different from their interaction in solution and may depend on the orientation of the complex and, in particular, on which part of the molecule binds to the surface. In the above schematic the adsorbed complex is represented in a generalized form (Me,L), without specifying a particular stereochemical arrangement. Furthermore the adsorbed complex may be formed by adsorption of a soluble complex or sequential adsorption of metal and ligand ions. Throughout this paper the phrase “adsorption of the complex” is meant to imply a reaction or series of reactions resulting in a single metal-ligand-surface species and does not refer to a specific mechanism by which such a species is formed. There are a t least three possible ways in which adsorption of a metal-ligand complex may depend on pH: (1)Adsorption of the complexed metal may be chemically analogous to and have similar pH dependence as that of the uncomplexed metal. (2) Adsorption of the complex may be similar to that of the free ligand. (3) The complex may not adsorb at all. These are the extreme, and therefore simplest cases. Real systems may be represented by combinations of the limiting cases, and adsorptive behavior of a given complex in a given system will depend on the identity of the surface, electrolyte composition of the solution, surface electrical potential, etc. In the first two situations, the complex may adsorb more or less strongly than uncomplexed species. Thus case 3 is a limiting case in which the complex adsorbs much less strongly than either free metal or free ligand. Adsorption of the complex in cases 1and 2 can be classified further on the basis of the strength of the complex-surface interactions as in Table I. When the surface is treated as analogous to a dissolved ligand, and surface reactions as analogous to hydrolysis reactions, the expected adsorptionpH patterns for the complexes are those in Figure 1. In the following section, chemical equilibrium expressions consistent with this conceptual model are presented and used to explore the behavior of the system as total metal, ligand, and H+ concentration vary. Chemical Reactions and Equilibria i n M e - L - s S y s t e m s
Mathematical Development. The total adsorption of metal in a system is the sum of free and complexed adsorbed species. Consider the case where adsorption of the complex
(4)
Me a
+ L 2 MeL Metallike systems require eq 1, 2, 4, and 5.Ligandlike systems require eq
1 and 3-5.
I---
Low adsorbate concentration Higher concentrotton
I
I
-
Low adsorbate Concentration I
PH
Flgure 1. Typical adsorption edges for metals, ligands, and complexes.
Dependence of fractional adsorption on total adsorbate concentration is discussed later in this paper. is similar to that of the free metal. A set of adsorption reactions and chemical equilibrium expressions describing all possible interactions in such a system is presented in Table 11. In the table, reactions 1-4 represent adsorption of free metal, a metallike complex, a ligandlike complex, and the free ligand, respectively. Reaction 5 represents formation of soluble MeL complexes. Since the sites for metal adsorption (ga)need not be the same as those for ligand adsorption (ge),a surface may be saturated with ligand molecules and still have excess metalsorbing sites available. Charge assignment to surface species is omitted in the equations for simplicity. The term “EDL” is included to account for the effects of the nonzero electrical potential near the surface. These effects constitute the major difference between complexation reactions in solution and adsorption (“surface complexation”) reactions. For the equilibria as written, the larger the EDL term, the greater is the electrical force opposing adsorption. Available adsorption models (12-15) estimate the EDL term, in various ways, and its value is not agreed upon. Since its exact value is not critical to this discussion, we choose to represent it in Volume 15, Number 9, September 1981 1051
a nonquantitative way rather than as a specific mathematical expression. Activities of bulk electrolyte at the surface and in solution are assumed to be constant. Since all of the equilibria are coupled, adsorption may affect speciation in solution in addition to altering total dissolved concentrations of metal and ligand directly. For instance, at low pH, adsorption of the ligand shifts the dissolved MeL/Me equilibrium toward the free metal. The adsorption reactions are written for the case that all protons released during the adsorption process originate at the surface. In reality, protons may also be released by acid/ base reactions of adsorbate metal and ligand ions. Such a process would not alter any of the qualitative conclusions presented here but would only change the numerical value of the adsorption equilibrium constants computed. The equations also include the implicit simplifying assumptions that only two metal species (free metal and one complex) need be considered and that each can be treated as if it adsorbs onto only one type of site by a single mechanism. The validity of these assumptions and the expected results if they are violated are discussed later. Manipulating eq 1,2, and 5 (Table 11)to get expressions for (-)/(MeL) and (SbMeL)/(MeL) and then adding these, we get
(SBMe) + (SbMeL)(MeL)
Kl(S"H,) K2(SbHb) Ks(H+)"(L)(EDL~) (H+Ib(EDLb) +
(6)
+
+
From eq 5 , (MeL)/{(Me) (MeL)) = Kb(L)/{l K5(L)). Multiplying both sides of eq 6 by this quantity and expanding:
(SBMe) + (SbMeL)-(MeL) + (Me)
Ki(SaH,)
[l
t
+ KS(L)](H+)~(EDL,)
The left-hand side of eq 6 is the ratio of adsorbed to dissolved metal, which we define as f . The right-hand side consists of two terms. The first ( F 1 )represents the contribution from adsorption of the free metal, and the second ( F z ) represents that from adsorption of the complex. The overall fractional metal adsorption, defined as (total metal adsorbed)/(total metal in the system), is f / ( f 1). Equation 7 applies for systems in which adsorption of the complex is metallike, i.e., adsorption of the complex is described by eq 2. If complexes adsorb analogous to free ligands,
+
eq 3, which describes adsorption of ligandlike complexes, replaces eq 2 in the set of simultaneous equations needed to characterize metal adsorption. The equations for adsorption of free metal (eq l),adsorption of free ligand (eq 4), and formation of soluble MeL complexes (eq 5 ) are applicable in either system. The ratio of adsorbed metal to dissolved metal for systems with ligand-like complexes is
~3K5(L)(SCHc)(H+)d(8) [I + Kb(L)](EDLc) f = F3
+ F4
(84
Model Predictions: pH-Adsorption Curves and Surface Speciation
In this section, the equations derived above are evaluated for several limiting cases. Expected changes in overall metal adsorption and surface speciation as functions of pH and ligand concentration are described for metallike and ligandlike complexes. Metallike Complexes. Overall fractional metal adsorption ( f / ( f 1))is outlined in Table IIIA for four limiting cases in which adsorption of the complex is metallike. The importance of ligand concentration in determining f changes dramatically among the four cases. In cases 1 and 4 a single species dominates both dissolved and adsorbed phases, and fractional metal adsorption is independent of ligand concentration. In case 2 the metal-ligand complex adsorbs less strongly than free metal, so increasing ligand concentration decreases overall metal sorption. Case 3 is the opposite situation in which the complex is strongly adsorbed, so increasing the ligand concentration increases metal adsorption. In these systems, the surface speciation, i.e., the relative concentrations of free and complexed metal species at the surface, is governed by the following expression:
+
If the complex adsorbs in a metallike fashion, it is likely that it binds to the same group of sites as the free metal. When this is the case, or when there is no s i t e w t a t i o n for adsorption of free metal or complex, the ratio ( S a H a ) / ( W )is constant. An additional simplification can be made for those systems in which pH-adsorption edges for the metal in the presence and the absence of ligand are approximately parallel, as is often reported. In such systems (H+)"(EDL,) = (H+lb(EDLb), and eq 17 reduces to
Table 111. Equations Describing Partitioning of Metal Ions in Systems Containing Complexing Ligands for Several Limiting Cases case
dominant dissolved metal specles
dom I na nt adsorbed metal specles
f
eq no.
A . Limiting Cases for Metallike Complexes
1 2 3 4
MeL (K5(L) >> 1) MeL Me (K5(L) F2) MeL Me
K2(Hf)-b(w)(EDLb)-1 K1K5-'(H+)-"(S"H,)(L)- (EDL,)K2K5(H+)-b(sbH,)( L)(EDLb)K~(H+)-"(S"H,)(EDL~)-
MeL Me MeL Me
K3(s"H,)(H+)d(EDLc)-1 K1K5- '(H+)-"(S"H,)(L)- '(EDL,)-' K~KS(H+)~(~"H,)(L)(EDL~)K~(H+)-"(S"H,)(EDL,)-
(9) (10) (1 1)
(12)
(13) (14) (15) (16)
where
Equation 18 indicates that surface speciation depends only on the free ligand concentration, and not on pH. If the free ligand concentration is fixed in a system, the same metal species is dominant a t the surface across the entire pH range. Adsorption of the ligand removes it from solution. In systems that are not site-limited for free ligand adsorption, dissolved ligand concentration decreases significantly at low pH. In such cases surface speciation varies with pH, with adsorption of the free metal favored at lower pH and adsorption of the complex favored at high pH. In many systems, however, ligands are present in great excess and total soluble ligand concentration is approximately invariant (e.g., C1- and S042in seawater). In such systems surface metal speciation is independent of pH. If the number of protons released when a complex adsorbs is different from that when the free metal adsorbs (a # b) or if EDL, # EDLb, then the slope of the pH adsorption edges in the systems with and without ligand will not be parallel. The slope of the adsorption edge will gradually change from that described by eq 12 to that described by eq 9 as ligand is added to the system. However, in all cases the exponent on the (H+) term is negative, so the general pattern of increasing adsorption with increasing pH will apply in any system. Ligandlike Complexes. Overall adsorption patterns for metals which form ligandlike complexes can be quite complex. If the complex adsorbs analogous to the free ligand, the ratio of adsorbed free metal to adsorbed complex is
As before, if all surface site types are available in excess, or if free and complexed metal bind to the same group of sites, the ratio (SaHa)/(SCHc)is constant. Equation 19 can then be rewritten:
where
The term (EDL,)/(EDL,) represents the ratio of electrical forces opposing adsorption of the ligandlike complex to that opposing adsorption of free metal. It is expected to increase with increasing pH, as does the term l/(H+)a+d.Equation 20 indicates that, in contrast to the case for metallike complex adsorption, the surface speciation is highly pH-dependent even at constant dissolved ligand Concentration. The complex dominates metal adsorption at low pH, and the free metal is dominant a t high pH. Plots of fractional metal adsorption vs. pH for these systems have four regions as illustrated in Figure 2. Region I: At low pH, surface speciation is dominated by strongly sorbing complex species, and nearly all of the metal is removed from solution, if enough sites are available. Region 11: At intermediate pH, the driving force for forming the surface-complex bond, and hence fractional adsorption of the complex, decreases with increasing pH.
U
I
W
-
. . . . . . ..I
m
U 0 v)
n U
U IW
E
......
I
c
I
I
nu
U
tJ'
'
E r n
u
W
a
PH Figure 2. Metal adsorption patterns expected in systems with ligandlike complexes. Each figure represents overall fractional metal adsorption in a system with fixed adsorbent concentration, fixed total metal concentration, and several different ligand concentrations. The strength of the bond between the surface and the ligandlike complex increases from A to B to C. In each figure, all of the curves intercept at a single pH, designated pH,.
Region 111: Also at intermediate pH, the driving force for forming the surface-free metal bond increases, and fractional adsorption of the free metal increases with increasing pH. Region IV: At still higher pH, adsorption of the free metal is so strong that nearly all of the dissolved metal adsorbs, assuming once again that sufficient sites are available. For a given ligand concentration, if the regions of decreasing complexed metal adsorption (region 11) and increasing free metal adsorption (region 111) are widely separated, the net adsorption curve is similar to Figure 2A. If the pH range of region I1 is more alkaline than that of region 111,regions I and IV overlap and the net curve is similar to Figure 2C. Figure 2B represents the intermediate case, where regions I1 and I11 overlap. I t is interesting to note that increasing adsorbent concentration shifts the adsorption edge for anions to higher pH and that for cations to lower pH, other solution conditions being constant. Thus, the degree of overlap of the edges is adjustable, and a system tends to shift from that represented by A to B to C as adsorbent concentration increases. Some important conclusions can be drawn from this qualitative analysis. If metal complex adsorption is analogous to that of the free ligand: (1) Assuming sufficient sites are available, fractional metal adsorption is near 100%at low and high pH and decreases to a minimum a t some intermediate pH. In some cases the decrease a t the minimum may be undetectably small. (2) There exists a value of pH, which we designate pH,, at which fractional adsorption of free metal equals that of complexed metal; at this pH overall fractional Volume 15, Number 9, September 1981
1053
metal adsorption is independent of ligand concentration. This can be shown from a mass balance on adsorbed metal species: adsorbed complex adsorbed free metal = total adsorbed metal
+
X
system
\com$exes
+
I
total free
fractional
( )( )
system
metal
overall fractional metal adsorption
in the system
Defining fractional adsorption of complexes (fcpx),free metal (fMe), and total metal ( f ~as )
(SLMe)
fcpx E
(w) + (MeL) .
fMe=
fT E
(SMe)
(SMe) + (Me)
(SLMe) t (SMe) (e) + (m)+ + (MeL)
(Me)
one can rewrite the above equality:
-
{(MeL)t (SCHc+dLMe)~fcpx + {(Me)+ (S*Me))fMe= (MeT)fT (21) Under conditions where fcpx = fMe KMeL) + (ScH,+dLMe) + (Me) + (SBMe))fcPx= (MeT)fT (22) Since the expression in brackets equals (MeT) fcpx = f M e
= fT
(23)
That is, the overall fractional adsorption is constant and equal to fMe and to fcpxand is independent of ligand concentration. Under conditions where fcpxis invariant with ligand concentration (dilute solutions), f~ a t pH = pH, will also be independent of ligand concentration. As ligand concentration increases in such systems, total metal adsorption increases at pH < pH,, decreases a t pH > pH,, and does not change at pH = pH,. Thus, a t pH near pH, a plot of fractional adsorption vs. p H has decreasing slope with increasing ligand concentration (Figure 2B). The case where fcpxdoes vary with ligand cohcentration is discussed below. Reevaluation of t h e A s s u m p t i o n s of the Model
At this point it is instructive to reconsider some of the assumptions of the derivation. The assumptions are as follows: (1)Only two adsorbing metal species need be considered. That is, free metal and a single complexed metal, species are the only adsorbing species. Alternatively, several complexes of a given metal-ligand pair may adsorb, but adsorption behavior of all complexes is equivalent. (2) Fractional adsorption of a species (e.g., fcpxor f ~ is a~ function of pH but not of the species concentration. 1054
Environmental Science & Technology
(3) Each species binds by a single mechanism, i.e., complexes are either “metallike” or “ligandlike”, and not both. Assumption 1. Most ligands form multiple complexes of the form MeL, with a given metal, where x varies between 1 and the coordination number of the metal ion. As the free ligand concentration in solution increases, the extent of complexation also increases. The model for adsorption of complexes has been developed by assuming that the adsorptive behavior of all complexes with a given ligand ( x > 0) is identical. While this assumption is an oversimplification, for most metal-ligand pairs the average value of x changes by one unit for a change in free ligand concentration of 1.5-3 orders of magnitude. In most experimental systems, and in natural systems other than estuaries or rivers in arid regions, ligand concentrations do not vary by more than 1order of magnitude. Thus, usually only one or two complexed metal species are of interest for a given ligand. The case of one dominant complex species has been modeled. The model could be extended to two complex species by treating the cases where the second MeL, complex adsorbs (i) analogous to the free metal or the free ligand, (ii) more or less strongly than the free metal, and (iii) more or less strongly than the first complex. Clearly the number of situations to consider grows rapidly as the number of adsorbing species grows. However, there is a drastic change in adsorption behavior only if the second complex adsorbs in a manner that is qualitatively different from the first (Le., free-metal-like vs. free-ligand-like). There are no experimental observations which necessitate such an explanation. Assumption 2. The assumption that fractional adsorption of a species is independent of its concentration will be violated (i) if the surface is composed of nonequivalent groups of sites, (ii) if availability of surface sites is a limiting factor, or (iii) if adsorption significantly alters the electrical interactions between the surface and adsorbing ions. (i) There is considerable evidence suggesting that oxide surfaces are composed of several distinct types of sites to which metal ions may bind (17,18).In such systems fractional metal adsorption at constant pH decreases with increasing total metal concentration, even when sites are available in excess and the EDL terms are approximately constant (Figure 3). By contrast, most of the available adsorption data for ligands can be explained without invoking surface heterogeneity (19,ZO). Extending these observations to adsorption of complexes, fractional adsorption of metallike complexes but not ligandlike complexes is expected to decrease with increasing concentration of complex (Figure 1).If evidence for adsorption of ligands on nonequivalent sites becomes available, the following discussion of metal adsorption on nonequivalent sites can easily be extended to include this case. Consider first a complex that has free-metal-like adsorptive behavior. As total metal concentration increases, fractional adsorption of both adsorbing species (free metal and complex) decreases at constant pH, and the overall fractional metal adsorption curve shifts to higher pH. If the ligand concentration is increased at constant pH while [ M ~ Tis] constant, the free metal concentration decreases and the complexed metal concentration increases. The fractional adsorption of free metal increases, and that of complexed metal decreases. Depending on which species is dominant at the surface, total metal adsorption may increase or decrease slightly relative to that expected for the case of equivalent surface sites. In the relatively few cases where these phenomena have been studied, the effect of surface site inhomogeneity has been much smaller than that due to the change in dissolved speciation and the difference in binding constants of the adsorbing species (16,18).Thus in systems with metallike complexes, surface site inhomogeneity may shift the fractional metal adsorption ) curve slightly, but probably will not significantly change its shape or its response to changes in other variables.
I
(Me) decreases from (l)t0(2)td3)
1
Figure 3. Effect on pH, of a shift in pH-adsorption edge for the metal ion. The lines labeled "Me" represent fractional adsorption of free metal. The shift from curve 1 to curve 3 could be caused by increasing ligand concentration or decreasing total metal concentration,either of which decreases free metal concentration. If ligandlike complexes are formed in a system with a nonuniform adsorbent surface, expected patterns for overall metal adsorption can be generated on the basis of the adsorption edges for the individual adsorbing species (Figure 3). As free metal concentration is decreased either by decreasing total dissolved metal concentration or by increasing ligand concentration, the curve for fractional adsorption of free metal shifts to lower pH, but that for the complex is not affected. The intersection of the free metal and complexed metal adsorption curves shifts to lower pH and greater fractional adsorption. As a result, in systems where the surface has several types of sites, curves for total fractional metal adsorption do not intersect at a single point, and there is a region where total metal adsorption is not a monotonic function of ligand concentration (Figure 4). For instance, at the p H labeled pH*, total metal adsorption increases as ligand concentration is increased from L1 to L2 and then decreases as ligand concentration is further increased to LB.As with metallike complexes, it is unlikely that surface site inhomogeneity significantly alters the general pH-adsorption patterns for systems with ligandlike complexes. Summarizing, surface site inhomogeneity alters the quantitative aspects of the model somewhat, but most quailitative predictions are unaffected. For systems with ligandlike
Figure 4. Expected changes in overall metal adsorption if the complexes are ligandlike and adsorbent surface sites are not all equivalent.
complexes, there exists a pH at which fractional metal adsorption is independent of ligand concentration only if metal-surface bond energies are approximately constant from site to site. Adsorption on a surface with nonuniform sites may be characterized by a pH region where metal adsorption is not a monotonic function of ligand concentration. For metallike complexes, surface site inhomogeneity may shift the curve of fractional metal adsorption vs. pH to higher pH but does not significantly alter its shape. (ii) The assumption that total available sites are not limiting is most likely to be violated at low pH in systems with large concentrations of anionic ligands. In essentially all natural aquatic systems, the total metal adsorption density is much less than the total number of available adsorption sites. Davis and Leckie (6, 20) have shown that in several systems the adsorption density required for site limitation is much less for anion adsorption than for cation adsorption. They attribute the difference to coverage of up to 20 sites per ligand molecule and of only one site per metal ion. In systems such as this the availability of sites may limit adsorption of ligandlike complexes as well as that of free ligands. The increase in metal adsorption at low pH would then be suppressed somewhat, and in extreme cases may not be evident at all. Other than this, site limitation does not alter the predicted adsorption behavior. (iii) If a charged species is specifically adsorbed, the electrical properties of the solid surface change in a way that opposes further adsorption. In most natural systems, trace-metal
Table IV. Analysis of the Assumptions of the Proposed Model for Adsorption of Complexes assumption
(1) all MeL, complexes
adsorb identically (2) fractional sorption
most llkely devlatlon
MeL,, MeL2,.. . have different binding strengths to surface surface site nonhomogeneity
of a species independent of species concentration surface sites limiting shifting IEP with ligand concentration
(3)complexes sorb by only
one mechanism
complexes can sorb in either metallike or ligandlike fashion
affects on predlcted sorptlon patterns
analysis more complicated and quantitative predictions altered some; qualitative pattern probably not altered much for ligandlike complexes, pH, decreases with decreasing free metal concentration;region of nonmonotonic relation between fractional metal sorption and ligand concentration appears: for metallike complexes, adsorption edge shifts to higher pH with increasing metal concentration not likely to occur for metallike complexes: decreases maximum sorption at low pH for ligandlike complexes negligible effect on metallike complexes: for ligandlike complexes, pH, decreases with increasing ligand concentration: metal sorption decreases some at low pH, but eventually attains 100% removed; region of nonmonotonic relation between metal sorption and ligand concentration appears analysis involves combining analyses for two limiting cases; no major effect
Volume 15, Number 9, September 1981 1055
1
I
PH
Figure 6. pH-adsorption edge for a complex which can adsorb in either
a ligandlike or metallike fashion.
Figure 5. Effect of total adsorbate concentration on the pH-adsorption edge for ligands which significantly shift the IEP of the adsorbent.Site
limitation causes the decrease in maximum percent adsorbed for ligand concentrations 2L4. concentrations are so low that their effect on the surface charge characteristics of the solid is probably negligible. However typical anionic ligand concentrations may significantly alter surface charge. For instance, Anderson et al. ( 2 1 ) found the IEP of amorphous aluminum hydroxide drops from mol/m2 to pH 8.5 to pH 5.0 as r A s T increases from 8.9 X 3.2 X mol/m2 (assuming 500 m2/g am-A1(OH)3). Pierce and Moore (22) also reported that the pH of the zero point of charge for amorphous from oxyhydroxide decreases as arsenite is added to the system. In such systems, the pH-adsorption edge for the ligand may shift to lower pH with increasing ligand concentration (Figure 5 ) . This situation can be treated in a way exactly analogous to the decrease in fractional adsorption of metal with increasing metal concentration. If the fractional adsorption curve for the complex shifts significantly, fractional adsorption of the complex decreases with increasing ligand concentration at low pH. In addition, the pH of minimum adsorption decreases, and there is a region where fractional metal adsorption is not a monotonic function of ligand concentration. Qualitative adsorption patterns are not altered. Assumption 3. The assumption that complexes adsorb by a single mechanism allows limiting cases to be analyzed more easily. Thus, the adsorption of the complex may either be metallike or ligandlike. If a complex can adsorb by both mechanisms, the pH-adsorption edge would be as shown in Figure 6. In that case, total metal adsorption is the sum of the adsorption of the three surface species: free metal, ligandlike complex, and metalike complex adsorption. The assumption that complexes adsorb by a single mechanism allows the more general case to be broken down into simpler, limiting cases for analysis. If this assumption is invalid, the analysis would be somewhat more complicated but the model would not require any major revision. The assumptions of the model and their implications are summarized in Table IV. Summary
A model for the effects of ligands on metal adsorption has been developed. The assumptions used to derive it are reasonable, though somewhat oversimplified. The deviations from model behavior that would be expected if the assumptions are invalid have been discussed. The model can be summarized as follows. (1)Metal-ligand complexes have pH dependence analogous to either the free metal or the free ligand. (2) The magnitude of the adsorption equilibrium constant for the complex can be less than, equal to, or greater than that for the free metal or ligand. (The case of totally nonsorbing complexes is a limiting case in which K M ~=L0.) 1056
Environmental Science & Technology
(3) If adsorption of the complex is similar to that of free metal: (i) The percent adsorbed-pH curve may be approximately parallel to that for the ligand-free system. (ii) For a given ligand concentration, the ratio of adsorbed complex to adsorbed free metal is given by eq 17. In many systems this ratio is approximately independent of pH. (iii) Increasing ligand concentration a t a given pH has a monotonic effect on fractional adsorption. (4) If adsorption of the complex is analogous to that of free ligand: (i) The complex is the dominant surface species at low pH, and the free metal is dominant at high pH. (ii) There is a minimum in total fractional metal adsorption. Depending on solution conditions and surface-binding constants, the minimum may be anywhere from near zero to near 100%adsorbed. (iii) There is a pH (pH,) at which fractional adsorption of the complex and fractional adsorption of the free metal are equal. Near this pH, the slope of a plot of fractional metal adsorbed vs. pH decreases with increasing ligand concentration. At pH < pH, metal sorption increases and at pH > pH, metal sorption decreases as ligand concentration increases, if each species binds to a single type of site. If either adsorbate binds to several different types of sites, or if the adsorption significantly alters the surface potential of the solid, there is a region where total metal sorption is not a monotonic function of ligand concentration. (iv) At very low and very high pH, nearly 100% of the metal is removed from solution as complexed and free metal, respectively. Literature Cited (1) Leckie, J. 0.;James, R. 0. In “Aquatic Environmental Chemistry of Metals”; Rubin, A. J., Ed.; Ann Arbor Science: Ann Arbor, MI, 1974; Chapter 1. (2) Hohl, H.; Stumm, W. J . Colloid Interface Sci. 1976,55,281. (3) James, R. 0.;Stiglich, P. J.; Healy, T. W. Discuss. Faraday Soc. 1975.59, 142. (4) Anderson, B. J.; Jenne, E. A.; Chao, T. T. Geochim. Cosmochim. Acta 1973,37,611. ( 5 ) Vuceta, J. Ph.D. Thesis, California Institute of Technology, Pasadena, CA, 1976. (6) Davis, J. A.; Leckie, J. 0. Enuiron. Sci. Technol. 1978, 12, 1309. (7) Elliot, H. A.; Huang, C. P. J . Colloid Interface Sci. 1979, 70, 29. (8)Bourg, A. C.; Joss, S.; Schindler, P. W. Chimia 1979,33, 19. (9) MacNaughton, M. G.; James, R. 0.J . Colloid Interface Sci. 1974, 47, 431. (10) Bourg, A. C.; Schindler, P. W. Chimia 1978,32,166. (11) Inoue, Y.; Minemori, M. Enuiron. Sci. Technol. 1979,13,433. (12) Davis, J. A.; James, R. 0.;Leckie, J. 0.J. Colloid Interface Sci. 1978,67,90. (13) Stumm, W.; Huang, C. P.;Jenkins, S. Croat. Chem. Acta 1970, 4 3 , 223. (14) Bowden, J. W.; Posner, A. M.; Quirk, J. P. Aust. J . SoilRes. 1977, 15, 121. (15) Anderson, M. A.; Malotky, D. T. J . Colloid Interface Sci. 1979, 72,413. (16) Benjamin, M. M.; Leckie, J. 0. submitted for publication in
Enuiron. Sci. Technol. (17) Beniamin. M. M.; Leckie, J . 0. J . Colloid Interface Sci. 1981. 79,209. (18) Benjamin, M. M.; Leckie, J. 0. In “Contaminants and Sediments”; Baker, R. A., Ed.; American Chemical Society: New York, 1980; Vol. 2, Chapter 16.
(19) Davis, J. A. Ph.D. Thesis, Stanford University, Stanford, CA, 1977. (20) Davis, J. A.; Leckie, J. 0. J . Colloid Interface Sci. 1980, 74, 32. (21) Anderson, M. A.; Ferguson, J. F.; Gavis, J. J . Colloid Interface Sci. 1976,54, 391.
(22) Pierce, M. L.; Moore, C. B. Enuiron. Sci. Technol. 1980, 14, 214.
Receiued for review May 2,1980. Accepted March 9,1981.
Crystalline Components of Stack-Collected, Size-Fractionated Coal Fly Ash Lee D. Hansen,*t David Silberman, and Gerald L. Fisher* Laboratory for Energy-Related Health Research, University of California, Davis, California 95616
Results of quantitative determinations of quartz, mullite, and magnetic iron oxides (magnetite and y-FezO3) are reported for size-fractionated coal fly ash. The concentrations of these crystalline phases are found to decrease as particle size decreases. Results of chemical analysis of the magnetic phase indicate that it crystallized from molten silicates during ash formation. Introduction
As part of a program to characterize the fly ash which is emitted by coal-fired power plants, qualitative identification and quantitative estimation of the crystalline components of four size-fractionated and one unfractionated fly ash sample are reported. Although fly ash is mostly amorphous to X-rays, the presence of small amounts of quartz, hematite, mullite, gypsum, magnetite, and ferrite have been reported (1-3). However, quantitative determinations of these mineral phases have not been reported, nor have the crystalline phases been studied as a function of particle size. A knowledge of the Crystalline phases is of importance in the consideration of the potential health effects of inhaled particles. Because of the refractory nature of the quartz, mullite, and magnetite phases, these materials will have long residence times in the pulmonary region of the respiratory tract if they are deposited there ( 4 ) .Therefore, it is important to know the particle size distribution and concentrations of these materials in stack-collected coal fly ash. Furthermore, it is generally recognized that crystalline siliceous materials are more toxic than amorphous compounds of the same composition. Such particles are known to have significant effects on lung cells ( 5 )and appear to be important toxicants to the pulmonary macrophage, the primary effector cell for lung immunosurveillance. Magnetite may also be a hazard to health because of its ability t o occlude biologically active transition-metal ions such as Mn and Ni by isomorphous substitution in the spinel crystal lattice (2).Magnetite could thus act as a slow release carrier agent for toxic elements. For this reason we have performed analyses of the magnetic phase for those metals which are likely to be associated with the magnetic fraction of the ash. The crystalline phases are important in determining the physical and chemical properties of the ash. Data on the crystalline phases may be useful in developing methods for resource recovery from, utilization of, and disposal of the ash (2).The mechanisms of formation of the various crystalline + On leave from the Department of Chemistry and the Thermochemical Institute of Brigham Young University, Provo, UT 84602. Present address: Battelle Columbus Laboratories, Toxicology/ Pharmacology Section, 505 King Avenue, Columbus, OH 43201.
0013-936X/81/0915-1057$01.25/0
phases are also of interest since a knowledge of these mechanisms may lead to methods for altering the composition of the ash. E x p e r i m e n t a l Section
Materials. The ash samples used in this study have been described in detail in previous papers (6-11). The size-fractionated samples were obtained from the stack breeching, downstream from the electrostatic precipitator (ESP), of a large coal-fired power plant in the southwestern U S . which was burning low-sulfur, high-ash coal. A sample collected from the ESP hopper was also analyzed. Min-u-sil216,obtained from Whittaker, Clark and Daniels, Inc., South Plainfield, NJ, was used as an a-quartz standard. The X-ray diffraction (XRD) pattern obtained from this material is given in Table I. A mullite standard was prepared by crushing and grinding a high-temperature furnace liner tube, dissolving the amorphous material in 1%HF, and washing with concentrated HN03, concentrated HC1,0.25 M EDTA, concentrated NHdOH, and methanol according to a recently reported procedure (12).The mullite standard was analyzed by atomic absorption spectrometry (AAS) (8),which gave 30 f 3% SiOz, 64 ii 1%A1203, 2.2 f 0.1% CaO, 0.22 ic 0.02% MgO, 330 f 50 ppm NazO, and 630 f 23 ppm KzO. The standard sample thus appears to be 89 f 1%mullite. The XRD pattern of the mullite standard is given in Table I. The acids used for trace-element analyses were HF (Baker Analyzed Reagent, 48%) and HC1 (G. Frederick Smith, 6 M, redistilled). Equipment. The X-ray diffraction patterns were obtained with a Diano XRD-8000 X-ray diffractometer using Ni filtered Cu K a radiation. AAS analyses were done with a Perkin-Elmer Model 306 AA spectrophotometer. Procedures. Specimens for identification by X-ray diffraction were mounted on flat glass slides either (a) by slurrying the solid with HzO, spreading on the slides, and drying at room temperature or (b) by sprinkling an excess of the dry ash onto double-stick Scotch tape on the slide and then gently tapping to remove the excess. Diffraction patterns were collected at 1.6' (28)/min from 2 to 70'(20) with a strip chart recorder speed of 1.6 in./min. Positions of peaks in the diffraction pattern could be located with an accuracy of f0.05'( 20). Estimations of quartz and mullite were performed by comparison of diffraction peak heights at 4.26 and 5.39 A to those obtained from a series of standard quartz-mullite mixtures. These wavelengths were chosen because there is no overlap between these peaks and those of any other component. Fly ash samples and the standards were prepared by spreading 25 mg of dry solid into a 1-cm diameter circle on double-stick Scotch tape attached to a glass slide. The dry solids were pressed into place with a second glass slide and
@ 1981 American Chemical Society
Volume 15, Number 9, September 1981
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