water interface II. Surface properties of amorphous iron oxyhydroxide and adsorption of metal ions

Surface Ionization and Complexation at the Oxide/Water Interface II. Surface Properties of Amorphous Iron Oxyhydroxide and Adsorption of Metal Ions J A M E S A. D A V I S 1 AND J A M E S O. L E C K I E Environmental Engineering and Science, Department of Civil Engineering, Stanford University, Stanford, California 94305 Received January 20, 1978; accepted May 3, 1978 The site-binding model for the electrical double layer of hydrous oxides reported in a previous paper is applied to the adsorption of metal ions from dilute solution and to complex heterogeneous systems, i.e., amorphous iron oxyhydroxide. More than one stoichiometric surface reaction is usually needed to describe the adsorption behavior of dilute heavy metal ions. If mass law equations for surface reactions of metal ions are corrected for effects of the electrostatic field at the interface, the calculated adsorption density depends upon the type of surface species formed. It is shown that calculations with surface reactions involving hydrolytic complexes of metal ions, e.g., Pb(II), Cd(II), Cu(II), Ag(I), are more consistent with experimental adsorption data than complexation by bidentate surface sites. A table of intrinsic surface complexation constants for various metal ions and oxide substrates is presented. Similar to results reported earlier for major electrolyte ions, the stability constants of surface complexes of heavy metal ions with the silica surface are significantly less than for other oxide surfaces. Empirical surface parameters for model calculations with iron oxyhydroxide are derived and the results are compared with experimental adsorption data for Cu(II) and Ag(I). INTRODUCTION

from swamping electrolyte solution, including cations, anions, and m e t a l - l i g a n d complexes. With increasing concern about trace elements as pollutants in recent years, there have been numerous studies o f adsorption o f metal ions by hydrous oxides (4-9). Many electrical double layer and adsorption models have been p r o p o s e d to account for the experimental observations. Some emphasize the importance o f the electrical double layer in controlling distribution o f solutes (10, 11), while others stress the specific chemical or coordinative interactions o f solutes with the oxide surface (5, 6, 12, 13). We have combined concepts from both approaches in the interfacial model presented here. The electrical double layer structure o f the site-binding model was developed for

In a p r e v i o u s paper (1), a site-binding model o f the electrical double layer o f hydrous oxide suspensions in aqueous solution was presented with an explanation o f the procedures for calculation o f surface parameters from experimental results. In combination with the solution equilibrium c o m p u t e r program, M I N E Q L (2), the model was used to describe the adsorption of ions and the development of surface charge and potential at the oxide/water interface in dilute and moderately concentrated simple electrolyte solutions. In this paper and a subsequent one (3), we use the site-binding model to describe the adsorption o f solutes 1 Presently at the Swiss Federal Institute of Technology (EAWAG), CH-8600 Diibendorf, Switzerland. 90 0021-9797/78/0671-0090502.00/0 Copyright © 1978 by Academic Press, Inc. All rights of reproduction in any form reserved.

Journal of Colloid and Interface Science, Vol. 67, No. 1, October 15, 1978

ADSORPTION

nonporous, crystalline oxides (1) and may not be the most appropriate for amorphous materials. Nonetheless, it will be shown that the experimental surface charge density of a complex heterogeneous system, i.e., amorphous iron oxyhydroxide, is also well described with two empirically estimated parameters, i.e., specific surface area and inner layer capacitance. It is important that we acquire the capability of modeling adsorption by amorphous iron oxyhydroxide because of its ubiquitous presence in natural systems and significant adsorptive properties (14). We first present resuits of surface characterization studies and then derive the surface parameters needed for adsorption modeling. MATERIALS

AND METHODS

Iron oxyhydroxide, Fe~O3" H20(am), was prepared in batch for each adsorption experiment with double-distilled water maintained at 25°C under a circulating purified nitrogen atmosphere. Sodium nitrate was added to achieve an ionic strength of 0.1 M after the addition of all reagents. Acidified stock ferric nitrate was then added and iron oxyhydroxide was precipitated by the dropwise addition of carbonate-free sodium hydroxide until a pH of 8.0 was attained. The precipitate was aged at pH - 8 for 4 hr before beginning the adsorption experiments. After this aging period, copper(II) or silver(I) was added from stock nitrate solutions. Since iron oxyhydroxide is not a stable solid phase, true equilibrium cannot be attained in short-term experiments. However, uptake of copper and silver was rapid and a quasi-equilibrium state was reached within a few hours. Equilibration periods of 2 and 4 hr were used for copper(II) and silver(I), respectively. Detailed results of the kinetic experiments are reported elsewhere (15, 16). The pH was measured using a Beckman Expandomatic meter No. SS-2 with a glass electrode and double-junction reference

OF METAL IONS

91

electrode with an outer salt bridge containing 10% KNOz. Copper concentrations were determined by flameless atomic absorption on a Perkin-Elmer Model 403 spectrometer with HGA carbon furnace. 11°~Ag was used as a radiotracer in Ag(I) adsorption experiments.

Characterization of Amorphous Iron Oxide Composition and morphology. Recent studies of iron (III) hydrolysis and precipitation (17, 18) provide information that is useful for characterization of the iron oxyhydroxide used in this study. Murphy et al. (17) characterized ferric hydroxy polycations formed in ferric nitrate solutions by electron microscopy and density gradient ultracentrifugation. The variables of interest were iron concentration, OH/Fe ratio, and aging time. In all solutions studied, the ferric polycations were spherical and in the same size range (15-30 A in diameter) after 3 - 4 hr of aging. A larger OH/Fe ratio generally increased the modal distribution of particle sizes for aging periods greater than a few days; however, after 3 - 4 hr of aging, there was little difference in particle size distribution as a function of OH/Fe ratio. Longer aging resulted in the formation of short rods consisting of two to five polycation spheres which subsequently formed raftlike structures composed of the rods. With further aging, the spheres comprising the rods became indistinct and coalesced. The individual rods and spheres gave no electron diffraction pattern, but goethite (a-FeOOH) was identified in rafts where coalescence occurred. Dousma and de Bruyn (18) presented evidence that above a critical pH (dependent on the solution conditions), iron(III) hydroxy polymers were formed that increased the optical density of the solution. These higher polymers probably correspond to the 15- to 30-/~-diameter polycation spheres observed by Murphy et al. (17), who also showed that a decreasing total iron concenJournal of Colloid and Interface Science, Vol. 67, No. 1, October 15, 1978

92

DAVIS AND LECKIE

tration lowers the average particle size produced. Since the iron concentrations used in our study were generally an order of magnitude lower than the lowest concentration (0.0165 M Fe) studied by Murphy et al., one can conclude that the diameter of the iron oxyhydroxide particles was approximately 20 A (or less) at the end of 4 hr of aging. Van der Giessen (19) also reported small particle sizes for an iron oxyhydroxide preparation, with the largest particles in the 20- to 30-A range. X-ray diffraction analysis of freeze-dried iron oxyhydroxide produced by our technique indicated that it was amorphous. However, the possibility of microcrystallinity undetectable by X-ray diffraction techniques exists (17, 19). Yates (20) determined the chemical composition of a similar amorphous iron oxyhydroxide as Fe2Oa'H~O. S p e c i f i c s u r f a c e area. BET analysis of nitrogen gas adsorption on freeze-dried iron oxyhydroxide after 48 hr of outgassing at room temperature indicated that the solid had a specific surface area of 182 m~/g. However, the BET surface area measurement must be considered with some caution. Yates (20) studied weight loss of amorphous iron oxyhydroxide as a function of time and temperature of outgassing and found that increasing amounts of chemisorbed water were released as the temperature was increased. Surface structure was probably altered by evolution of chemisorbed water. Yates presented evidence that outgassing of iron oxyhydroxide e v e n at r o o m t e m p e r a ture caused some surface decomposition. This may result in an underestimate of the actual specific surface area. Other reported values for specific surface area of iron oxyhydroxides prepared in a similar manner (measured by BET analysis) are 320, 257, and 159 m2/g (20-22). Specific surface area measurements can be made in situ by negative adsorption to avoid the problem of surface decomposition. The surface area of amorphous iron oxyhydroxide has been estimated in the

range of 270-335 m~/g by negative adsorption of Na + at pH 4 (15) and - 7 0 0 m2/g by negative adsorption of Mg 2+ at pH 5 (22). However, it has been shown that this method applies only to smooth, nonporous surfaces (23). Irregularities in surface structure and pores smaller than the double layer thickness are not detected. Thus, both negative adsorption and BET measurement techniques may result in low estimates of specific surface area for iron oxyhydroxide because of porosity and surface decomposition, respectively. Assuming spherical particles of 20-A diameter and the density (3.57 g/cm a) given by Murphy et al. (17), a specific surface area of 840 m~/g can be derived. The actual specific surface area may be less, but is probably larger than that estimated by negative adsorption (-300 m~/g). S u r f a c e c h a r g e a n d p H e z c . Surface charge density (tr0) on hydrous oxides is defined by the net uptake of protons by the surface and is determined by potentiometric titration (24), i.e.,

Journal of Colloid and Interface Science, Vol. 67, No. 1, October 15, 1978

O'o = F(FH+ - FoH-) = F ( C A - CB + [OH-] - [H+])/A,

[1]

where tr0 has the units coulombs per square centimeter, A is the total surface area in suspension, and CA and CB are the concentrations of acid and base after addition. We have determined pHvzc (pH at which tr0 is zero) for iron oxyhydroxide in NaNOa solution by a salt titration method (25). The titration was accomplished by observing the change in pH (ApH) upon the addition of salt to an iron oxyhydroxide suspension. The added salt shifts pH toward pHvzc by increasing [tr0[. At pHpzc, there should be no change in pH upon the addition of salt to the suspension. Figure 1 shows ApH as a function of pH measured at low ionic strength (1.5 x 10-aM NaNOa). A pHvzc of approximately 7.9 is indicated, which is in good agreement with the value of 8.0 reported by Yates (20) and others (26).

ADSORPTION OF METALIONS +

0.75

I

I

I

93

I

I

Fe T 5.6x 10"aM + 0.50

NaNn

•

1115x 10"2M

*0.25

-r o.

0.00

-0.25

-0.51 6

I

I

I

I

I

I

6.5

7

7.5

8

8.5

9

9.5

pH ( N o N 0 a , I.5 x I()aM) FIG. 1. Change in pH (ApH) upon the addition of sodium nitrate to amorphous iron oxyhydroxide suspensions as a function of pH measured in low ionic strength medium (1.5 x 10-3 M NaNO3).

pHvzc is -7.9. The net uptake of protons by the surface was determined by potentiometric titration of an iron oxyhydroxide suspension with carbonate-free 0.1 N NaOH and comparison with a titration of the electrolyte medium (0.1 M NaNO3). Since the pHazc is known, surface charge at any pH can be

,o

I I

DIFF SE LAYER

v,

0

,

0

d

F~. 2. Schematicrepresentationof the charge distributionat an idealizedplanarsurfaceand the potential decay away from the surface.

calculated from the difference between the two titrations. The surface charge data are presented in the next section with model calculations.

The Electrical Double Layer of Amorphous Iron Oxyhydroxide Surface ionization. A method was presented previously (1) for calculation of surface ionization constants to describe the surface equilibria o f crystalline oxides in simple electrolyte solutions. Intrinsic surface ionization constants, i.e. /¢'int /(int were calculated at low ionic strength and intrinsic surface complexation constants, e.g., *lgint~x Na +, *iF'int,. NO:, at moderately high ionic strengths (0.1-0.5 M) when surface complexation dominates the surface proton balance. The importance of surface complexation in simple electrolyte solutions has recently been confirmed experimentally (45, 46). In addition, estimates of intrinsic constants can now be improved by a new double extrapolation technique (27, 37). The mathematical formulation for surJournal of Colloid and Interface Science, Vol. 67, N o . 1, O c t o b e r 15, 1978

94

DAVIS AND LECKIE

face charge was developed with a double layer model of a planar, nonporous surface (Fig. 2). Although surface bonding and surface geometry may differ between amorphous iron oxyhydroxide and crystalline, nonporous oxides, one may nevertheless describe the development of surface charge on iron oxyhydroxide with the same mathematical formulations. Furthermore, the graphical technique for determination of intrinsic surface ionization constants does not require that the specific surface area of the solid be known. Figure 3 illustrates calculations of pKiant and e"*t"i"t,~ NO3- for iron oxyhydroxide in NaNO3 solution. The total site density, Ns, used to make the calculations in Fig. 3 was determined by rapid tritium exchange with the surface (20). Since the reported site density (11.4 sites/ nm 2) was dependent on the specific surface area determined by BET gas adsorption, a site density based on mass may be more appropriate for amorphous materials. In these calculations and elsewhere (3), I

1

we have used 9.85 × 10-3 moles sites per gram of Fe203-HzO(am), recalculated from the raw data of Yates (20). Inner layer capacitance and specific surface area. In addition to intrinsic ionization and surface complexation constants, the inner layer capacitance and specific surface area are needed to calculate surface charge density using the site-binding model. In the examples given previously (1), specific surface area was a known, experimentally determined quantity and the inner layer capacitance was used as a fitting parameter. However, there is a large range of estimates by different techniques for the specific surface area of amorphous iron oxyhydroxide, e.g., 157-840 mZ/g; hence, neither the specific surface area nor inner layer capacitance is well known. Nonetheless, using the estimates for acidity and surface complexation constants (Fig. 3), it is possible to select a combination of the two parameters that gives good agreement between calculated and experimental

I

I

r

~'K int

_.,.~P

,'~ o

I

Fe203" H20

zs-c

.o;

6

$o. ,

Int

jP~*, O° o.

5

4

3

0.00

I

I

I

I

I

[

0.01

0.02

0.03

0.04

0.05

0.06

ct

0.07

+ = ~./N=

FIG. 3. pQa, a n d P * QNO~- as a f u n c t i o n o f f r a c t i o n a l i o n i z a t i o n for a m o r p h o u s iron o x y h y d r o x i d e in N a N O 3 . Journal of Colloid and Interface Science, Vol. 67, N o . 1, O c t o b e r 15, 1978

95

ADSORPTION OF M E T A L IONS 70

I

I

I

0 ' % ~

6.5

I

2 X I0-3 M Fer 0.18 g / l

Fe20a • H=O

0.1 M NON03, 25*C ~

[ SOH]T

1.75 x IO-a M

6.0 pH 5.5

5.0 CI I00, 140 p.F/cm=

CI I00, 140p.F/cm z

400 m2/gm 4.5

0.25

0.50

600 rnZ/gm

I

I

I

[

0.7,5

1.00

1.25

1.50

1.75

¢re (mole/,t X I0 -4) F I 6 . 4 . Surface charge (in moles per liter) of amorphous iron oxyhydroxide in 0.1 M NaNOa as a function o f pH. Solid lines are model calculations with different values o f C1 and specific surface area. Circles are experimental data.

surface charge. Figure 4 shows calculated charge density, derived from experimental surface charge (solid lines) for different measurements of specific surface area and values of specific surface area and inner potentiometric titration, is overestimated layer capacitance in comparison to experi- due to errors associated with the techniques mental data (circles). While neither quantity for surface area measurement. The actual is estimated unambiguously by this method, specific surface area of amorphous iron 140 /.,F/cm2 (C1) and 600 m2/g are rea- oxyhydroxide may be greater than that sonable values that adequately describe measured by either BET gas adsorption or the experimental surface charge data. Al- negative adsorption. Figure 5 shows dethough the specific surface area (600 mZ/g) rived values of surface charge density as a may appear large, the negative adsorption function of ApH (pH - pHezc), using three experiments and the small particle size of different estimates for specific surface area the precipitate (17, 19) support the proba- from BET gas adsorption (182 m2/g), negability of a high specific surface area. Spe- tive adsorption (300 m2/g), and the empirical cific surface areas as large as 700 m2/g have estimate of Fig. 4 (600 m2/g). The solid lines been observed for other gelatinous oxide in Fig. 5 are experimental measurements of precipitates (28). surface charge density ofgoethite and hemaYates (20) observed a significantly larger tite (20). Thus, the surface charge density surface charge density for amorphous iron of amorphous iron oxyhydroxide is not sigoxyhyroxide in comparison to other oxides nificantly larger than that of crystalline (29, 30) and proposed that this phenomenon iron oxides if the specific surface area is was caused by the existence of a porous about 600 m2/g. double layer with very high capacitance. Amorphous iron oxyhydroxide is comHowever, it is also possible that the surface monly found in natural aqueous systems as Journal of CoUoid and Interface Science, Vol. 67, No. 1, October 15, 1978

96

DAVIS

50

AND

LECKIE

i

'0

182

A

40

0

i

i

mZ/gm 300 mZ/gm 600 m2/gm

0

Q

Hematite 30

A

0

Goethite

F e 2 0 3 • H20 0.1 M N o N O 3

( 'Votes [ 1975) . . . .

l 0.1 M KNO 3

A O

20 ,r,j :L

A

~o

A

0

I0

_..2L_ t

I

t

I

I

I

-3

-2

-I

-10

0 -20

--/

I

I

o

I

A

2

A p H : pH - pHpzc

FIG. 5. Surface charge densities of various iron oxides as a function of ApH. Points are derived from experimental measurements of the surface charge of amorphous iron oxyhydroxide and estimates for specific surface area. The solid lines are from the experimental data of Yates (20). a discrete mineral phase and as a surface coating on particulate matter (14). Surface binding by this material may control the distribution and accumulation o f many trace metals in natural waters (31). It is important that phenomenological adsorption models are developed that allow consideration o f such complex solids. Many questions remain concerning the surface structure of iron o x y h y d r o x i d e ; however, we feel that the surface parameters estimated empirically in Fig. 4 [140/xF/cm z, inner layer capacitance (C1); 600 m2/g, specific surface area] are adequate for our modeling purposes. While other combinations o f these two quantities might be found to describe

the surface charge data, such a combination would require that either C1 or specific surface area be greater than the present estimates. An increase in either quantity does not seem justified. The value estimated for C1 (140/zF/cm 2) is in good agreement with values found for several crystalline oxides (1), as shown in Table I.

Adsorption of Dilute Metal Ions Significant advances have been made in recent years in the development of phenomenological models to describe trace metal adsorption on hydrous oxides. In particular, the concepts o f surface ionization and com-

Journal of Colloid and Interface Science, Vol. 67, No. 1, October 15, 1978

ADSORPTION OF METAL IONS

97

TABLE I Parameters Used in Electrical Double Layer Modeling

Ref.

Surface area (m~/g)

~/-AI~Oa/NaCI

(38)

117

a-FeOOI-I/KNO3

(20)

a-FeOOH/NaC1 Fe2Oa' H20(am)/ NaNO3

System

Surface site density (sites/nm s) and Ref.

pKtm °~

pK~t

(39)

5.7

11.5

9.2(?)

7.9(?)

48

16.8 (20)

4.2

10.8

8.9

6.1

I00

(40)

32

16.8 (20)

4.9

a

a

6.6

a

This study

600(?)

0.00985 moles/g of Fe2Oa' H20 (20)

5.1

10.7

9.0

6.9

140

TiOz/LiNOa

(20)

20

12

(20)

2.6

9.0

7.1

4.5

140

TiOz/KNO3

(20)

20

12

(20)

2.6

9.0

7.1

4.5

100

TiO~/Mg(NOa)2

(20)

20

12

(20)

2.6

9.0

4.5

140

SiO2/KCI

(41)

170

5

(42)

a

7.2

ND

125

8

t*t p • K¢.u~

12.5 (MgOH +) 6.7

p*K~

C1 (/~F/em2)

100-120(?)

a Not determined.

plexation introduced by Stumm, Schindler, and their co-workers (5, 6, 12, 30, 32, 33) have aided in the understanding of surface chemistry and complex adsorption phenomena. We have used some of these concepts (1) to add a model for oxide surface ionization and complexation to the solution equilibrium computer program, MINEQL. 2 As a further development, we have introduced exponential terms to the mass law expressions to correct for effects of the electrostatic field on surface equilibria. While this physical refinement of the model may be less complete than other proposed models (10), it allows the consideration of adsorption of major electrolyte ions and dilute solutes simultaneously, and self-consistent calculations can be made that take into account the net interactions of all ions at the oxide/water interface. Formulation of surface equilibria. Surface reactions of cations in dilute solution may be written in a manner analogous to that of a swamping electrolyte cation (1). 2 A manual (2) describing the solution equilibrium computer program, MINEQL, is available from the Department of Civil Engineering, Massachusetts Institute of Technology, Cambridge, Mass. 02139.

For example, one may write the following reaction for Pb 2÷ with an oxide surface site, i.e., *K~t~+ SOH + Pbs2+q S O - - P b 2+ + Hs+,

[2]

where the subscript s denotes a surface concentration (1). The mass law equation is written in terms of bulk concentrations and the mean electrostatic potentials at the planes of mean charge, o-0 and o-8 (Fig. 2),

i.e., [SO-_Pb 2+] =

[SOH][Pb 2+] [H + ]

× exp[(e~0 - 2ed~a)/kT]*K~btZ+. [3] For purposes of mathematical simplicity, we shall assume that the mean location of all specifically adsorbed ions (including the swamping electrolyte ions) is the era plane, and that ~ is the mean electrostatic potential at that plane. Since Pb s+ has two units of charge, the electrostatic component in Eq. [3] includes the term [exp(-2e~b~/kT)]. The equations for o-o and tro are general

Journal of Colloid and Interface Science, Vol. 67, No. 1, October 15, 1978

98

DAVIS AND LECKIE

and must include all surface species. For example, in dilute Pb(II) solutions with NaNO3 swamping electrolyte, the following equations define o-0 and tra, i.e.,

SOH + Pb~2+ + HsO~

~ro = [SOH2 +] + [SOHs+-NO3 -]

(SOH)2 + Pb~s+ + HsO ~

*/,-" i n t ,x PbOH +

[ S O - - P b O H +] + 2H~+,

[8]

- [ S O - ] - [ S O - - N a +] - [SO--PbS+], [4] 0-8 = [ S O - _ N a +] + 2 [ S O - - P b 2+]

[SOH2+-NO3-].

-

SO-\pbs+ SO-

The charge/potential relationships are written (1) as follows: [6]

qJe - Oe = - o ' a / C z .

[7]

The symbol/3 is used for the stability constant when a surface reaction involves a bidentate surface site, e.g., Eq. [9]. These additional surface complexes must be included in the equations for O'o and o-~ (Eqs. [4] and [5]) if they are considered in the calculation of system equilibrium.

Adsorption of metal ions may result in the formation of more than one type of surface complex. For example, to model adsorption of Pb(II) on an oxide surface, one may want to consider surface reactions other than Eq. [3], e.g., I00

[

0

I

Pb(ll) T

Y - AIzO 3 0.1 M Hohl

Example

System: Pb(II)/y-AIsOa

Figure 6 shows data of Hohl and Stumm (5) for Pb(II) adsorption on y-AlsOa in 0.1 M

I

I

f

2.9x 10-4M 1371 mZ/,I

O

NoClO 4 ond

Stumm

E,,.

÷

[,]

(1976)

6O

so-

O

j

O o

o -J

40

N ÷ N ' SO--Pb t+] + [SO-- PbOH+ ]

20

,~,~.,.-'

O E.

0 35

I 4.0

I 4.5

[9]

/

[5]

too - O~ = O'o/C,,

+ 2Hs +.

I 5.5

5.0

[,1, I 6.0

,÷] I 6.5

I 7.0

pH FIG. 6. Adsorption o f Pb(II) on T-AlsO3 in 0.1 M NaCIO4. Circles are experimental data o f Hohl and S t u m m (5). T h e solid and d a s h e d c u r v e s are model calculations that consider different surface reactions o f Pb(II). Journal of Colloid and Interface Science, Vol. 67, N o . 1, O c t o b e r 15, 1978

99

A D S O R P T I O N O F M E T A L IONS

NaCIO4. The surface properties of y-A12Os in NaCIO4 are given in Table I. One begins by considering the simplest equilibria possible, divalent lead ion, Pb 2+, reacting with one oxide surface site, Eq. [2]. With this single surface reaction the calculated Pb(II) adsorption as a function of pH is in poor agreement with the experimental data. The model prediction in this case is the lower dashed curve in Fig. 6, and the circles are experimental data of Hohl and Stumm (5). 3 An increase in *K~ ~+ does not improve the agreement, as it simply shifts the calculated adsorption density to a higher plateau without increasing the pH dependence. The slope of the experimental adsorption density/pH data suggests that more than one proton is released per sorbed Pb(II) ion. Indeed, Hohl and Stumm determined by potentiometric titration that the number of protons released per Pb(II) ion adsorbed varied between one and two and averaged about 1.5 in the pH region of 4-6.5. There are two additional surface reactions which may account for the release of two protons per sorbed Pb(II) ion. The first involves adsorption of the monohydroxolead(II) complex, PbOH +, Eq. [8]. The second possibility is the formation of bidentate surface complexes, Eq. [9], as proposed by Schindler et al. (6, 32) and used by Hohl and Stumm (5) to describe their experimental adsorption data. The primary difference between the two reaction schemes, Eqs. [8] and [9], is the origin of the second proton released. The additional proton released upon adsorption of a hydrolytic complex is from the hydration sheath of the adsorbed metal ion. This is thermodynamically equivalent to surface hydrolysis of metal ions adsorbed at the surface, as shown by MacNaughton and James (8). In the formation of a bidentate surface 3 Although the surface equilibria presented in Eqs. [2], 18], [9], and others are written in terms o f concentration, [H+], the model calculations and experimental measurements presented in all figures are in terms o f pH, i.e., - l o g activity o f H+(aq).

complex (Eq. [9]), however, the second proton is released by ionization of a second surface site. This is an extremely important difference in terms of the way the electrostatic component of the mass law equations is formulated. In adsorption of a hydroxo complex, one proton is released from the o-0 plane and one from the hydration sheath of Pb 2+, i.e., the o-8 or o-a planes (i.e., cis or trans position of the hydroxyl ligand). Assuming for the moment that the net charge in the o-8 plane is + 1, the mass law expression for Eq. [8] is written as follows: [SOH][Pb2+] [H+]2

[ S O - - P b O H +] =

× exp[(et~o- e~8)/kT]*Kip~o.+.

[I0]

If Pb2+ is adsorbed as a bidentate surface complex, Eq. [9], two protons are released from surface sites, the ~0 plane, and the mass law expression is written as follows:

[Pb 2+]

/ SO-

1 \pb2+ / J

SOH =

[H+]Z

SO× exp[(2etk0 - 2eqJa)/kT]*fli,nbt*+.

[11]

Although Eqs. [10] and [11] are both dependent on [H+] -2, the electrostatic components of the mass law expressions differ. In adsorption of a hydrolytic complex, the electrostatic term is exp[(eq,0- eqJs)/kT], since there is a single unit of charge in each plane. For formation of a bidentate surface complex, the electrostatic term is exp[(2etO0- 2etkD/kT], as there are two units of charge in each plane. Since qJ0 and qJ8 are dependent on [H +] (1), the pH dependences of Eqs. [10] and [11] are not identical. The results of model calculations assuming Eq. [2] alone are compared with results Journal of Colloid and Interface Science, Vol. 67, No. 1, October 15, 1978

100

DAVIS AND LECKIE

assuming Eq. [2] in combination with either Eq. [8] or Eq. [9] in Fig. 6. Significantly better agreement is achieved in the case of adsorption of a hydrolytic complex, the solid curve in Fig. 6. The total adsorption density of Pb(II) is a sum of the surface concentrations of [ S O - - P b 2+] and [ S O - PbOH+]. S O - - P b O H ÷ is the predominant Pb(II) surface species for pH >- 5. The adsorption of monohydroxo complexes is consistent with the frequent occurrence of adsorption edges near the pH at which hydrolysis of the free metal ion begins (7). The model calculations in the case of Pb(II) complexation by bidentate surface sites (upper dashed curve, Fig. 6) are not in good agreement with the experimental data. The calculated curve is relatively insensitive to pH, even though bidentate surface complexes are the principal Pb(II) surface species comprising the total adsorption density. This occurs because the quadratic dependence on pH in Eq. [11] is counterbalanced by an increased dependence on ~0. As the pH increases, the formation of dissociated surface sites, SO-, is favored by decreasing [H+]. However, this effect is offset in the mass law expression, Eq. [11], by a decrease in 00, reducing the coulombic repulsion of [H +] at the tr0 plane. Thus, the pH dependence of complex formation with bidentate surface sites is not significantly greater than that considered for one proton released per sorbed Pb(II) ion, Eq. [2]. Since bidentate surface complexes comprise the bulk of the total adsorption density in the calculation shown, a change in the value of ,fl~t~+ will not alter the pH dependence or improve the fit to experimental data. It should be emphasized at this point that the dissimilarity between our calculations and those of Hohl and Stumm (5) result from different models of electrical double layer structure. The effect of surface potential, 00, on the ionization of surface sites is greater in our model. Furthermore, Hohl and Stumm did not include electrostatic

components in the mass law expressions for Pb(II) adsorption. In addition, the experimentally observed proton displacement due to Pb(II) adsorption, determined by potentiometric titration, is not in agreement with their calculations. The authors suggested that the bi-dentate surface complex was necessary to explain the average of 1.5 protons released per sorbed Pb(II) ion. However, in an additional graph [Fig. 4b in Ref. (5)], the authors showed that h, the calculated average number of sites bound to each Pb(II) ion, was only slightly larger than one, suggesting that bidentate complexes formed to a very small extent in comparison to 1:1 complexes. Adsorption of a monohydroxolead(II) complex would seem to be a more consistent explanation. In addition to the better agreement between calculated and experimental adsorption density (Fig. 6) using Eqs. [2] and [8], the calculated proton displacement in this study is more consistent with the potentiometric titrations of Hohl and Stumm (5). The average number of protons released per sorbed Pb(II) ion varies as a function of pH. At a low pH, when S O - - P b 2+ comprises most of the total adsorption density, the average is near one. As pH increases and the formation of S O - - P b O H + becomes significant in the surface speciation of Pb(II), the average increases to nearly two protons per Pb(II) sorbed. Although the calculated averago number of protons released varies more as a function of pH than the experimental data reported (5), it is more consistent with the stated average of 1.5 protons released per Pb(II) sorbed. DISCUSSION

General Applicability Using the same assumptions as above, we find similar results in applying the site-binding model to other studies of cation adsorption on oxide surfaces. Figures 7 - 9 demonstrate the agreement between calculated adsorption density and experimental data for

Journal of Colloid and Interface Science, Vol. 67, No. 1, October 15, 1978

ADSORPTION OF METALIONS I

0

Cd(Tr) T

TiO 2

80

I

i

101 I

I

2XlO -4 M

2 0 0 m2/_/

0.01 M KNO 3 , 25"C Stiglich (1976) "o (p

60

.od,I C,lc.l,,io.

IO 'V

E

._=

E

"0 o 0

/

[ SO--,CdOH ÷] O [SO--Cd=']

0

40

20

0

4

4.5

,5

5.5

6

6.5

7

pH

FIG. 7. Experimentaldata of Stiglich (34) and model calculationfor Cd(II) adsorption on TiO2 in 0.01 M KNOa. Adsorptiondensityis the sum of the surface concentrationsof [SO--CdOH+] and [SO--Cd2+]. Cd(II) adsorption on TiO2 (7, 34) and Cu(II) and Ag(I) adsorption on amorphous iron oxyhydroxide (15). In Figs. 7 and 8, the calculated curves for Cd(II) and Cu(II) uptake simulate the experimental data more successfully when surface reactions are written in terms of adsorption of a monohydroxometal(II) complex, rather than complex formation with bidentate surface sites. The surface parameters of the solid substrates and intrinsic surface complexation constants for metal ions are presented in Tables I and II, respectively. Silver(I) is unique in comparison to other transition metal ions because of the monovalent charge of the aquo ion and its weak hydrolysis behavior. These properties may influence adsorption, since we find that uptake of AgO) is more dependent on the

amount of solid present and somewhat less dependent on the pH than observed for many divalent transition metal ions (15, 16). Nonetheless, calculations with the sitebinding model indicate that the predominant surface species formed is the hydrolyzed Ag(I) ion, similar to results found for divalent metal ions. The dashed curve in Fig. 9 considers only the coordination of Ag ÷ by an oxide surface site, i.e., ./d-lnt+ i x Ag

SOH + Ags+4 S O - - A g + + He+.

[12]

Agreement between this calculated curve and the experimental adsorption data is unsatisfactory. The pH dependence of calculated adsorption density can be increased by considering an additional surface reacJournal of CoUoid and Interface Science, Vol. 67, No. 1, October 15, 1978

102

DAVIS AND LECKIE I00

I

0 Cu(I'I')T

I

I

0

10-6 M

(33

Fe T 10- 3 M 80

"0

0

0.1 M NoNO 3 , 2 5 * C

60

0 0 'I0

Q. CL 0 (.J

40

Model

[

Results

[so--cuo, ÷]

°

20

0 I

4.0

I

I

I

1

I

I

4.5

5.0

5.5

6.0

6.5

7.0

pH

FIG. 8. Adsorption of Cu(II) on 10-3 M amorphous iron oxyhydroxide in 0.1 M NaNOa. Circles are experimental data and the solid line is a model calculation considering Cu2+ and CuOH ÷ surface equilibria. tion involving the m o n o h y d r o x o complex, AgOH°(aq) (solid curve, Fig. 9). Thus, it would a p p e a r that adsorption of Ag(I), as well as divalent cations, involves formation o f surface c o m p l e x e s with hydrolyzed metal ions. T h e s e c o m p l e x e s can be considered as mixed ligand c o m p l e x e s in the sense that O H - and an oxide surface site each contribute to coordination o f the metal ion. Careful studies of the kinetics of surface reactions will be n e c e s s a r y to determine w h e t h e r hydrolysis occurs first in solution or at the surface following adsorption of an aquo metal ion.

Predictive Capability Similar to other p r o p o s e d adsorption models, the site-binding model presented here can be used in a predictive m o d e if

estimates of the stoichiometry o f surface complexation reactions and intrinsic constants are available f r o m other studies. F o r m a n y solutes, the chemical c o m p o n e n t of the bond energy with an oxide surface site m a y be a p p r o x i m a t e l y the same f r o m one oxide to another. Accordingly, intrinsic surface complexation constants k n o w n from results of adsorption experiments with one oxide can be used as a first estimate for the stability constants with a n o t h e r oxide. The intrinsic surface c o m p l e x a t i o n constants in Table II cannot be used directly for predictive purposes because the surface reactions are written in terms of nonionized surface sites, [SOH], and hence are dependent on the intrinsic acidity constants o f the particular oxide studied. H o w e v e r , surface reactions can also be written in terms of

Journal o f CoUoid and Interface Science, Vol. 67, No. 1, October 15, 1978

A D S O R P T I O N O F M E T A L IONS

I00

I

I

I

I

4XlO-7 M /I~T IO_SMF.eT

/ //~0

/

0.1 M NoNOa, 25"C

80

103

0 ~,

sttts S

,,' "o 0 J=l

60

it int p KA¢+ 4.9

m

iI

_~ulnt ,,.,, P r'AgOH ' r ' l

("~/

0;,"

40

/I t s"

I

11

[ SO-" Ag+ ] .4tvint~

A©

PKAg+ 4.6

ss J

20 --"

0 4

5

6

7

8

9

pH FIG. 9. Model calculations and experimental data for Ag(I) adsorption on 10-3 M a m o r p h o u s iron o x y h y d r o x i d e in 0.1 M NaNO3. D a s h e d curve considers only Ag + equilibria; p*KAs+,tnt4.6. Solid c u r v e considers both Ag + and A g O H ° (aq) surface reactions; p*KAg+, 4 . 9 , 1 n t " p*KAgoH,lnt12.1.

ionized sites, i.e., SO- or SOH2 +. The former type of reaction is designated by an asterisk superscript in the stability constant ( p * K i n t ) , while the latter type is simply p K lnt

(1). When written in terms of charged sites, the intrinsic constants are not dependent on the acidity constant of the surface and are directly comparable.

T A B L E II Intrinsic Surface C o m p l e x a t i o n C o n s t a n t s for Cation Adsorption on H y d r o u s Oxides

Figure

System and Ref.

Adsorbing species

p*K lnt

Surface reaction

6

Pb(II)/T-AI203/0.1 M NaCIO4 (5)

Pb 2+ PbOH ÷

5.0 10.3

S O H + pb 2+ m [ S O - - P b *+] + H ÷ S O H + Pb 2+ m [ S O - - P b O H +] + 2H +

7

Cd(II)/TiO2/0.01 M KNO3 (7, 34)

Cd 2+ CdOH +

1.8 8.7

S O H + Cd *+ m [ S O - - C d 2+] + H + S O H + Cd 2+ m [ S O - - C d O H +] + 2H +

8

Cu(II)/Fe2Os'H20(am)/0.1 M NaNOa This s t u d y

Cu 2+ CuOH +

4.1 9.0

S O H + Cu ~+ ~ [ S O - - C u ~+] + H + S O H + Cu 2+ ~ [ S O - - C u O H +] + 2H +

9

Ag(I)/Fe203' H~O(am)/0.1 M NaNO3 This s t u d y

Ag + AgOH °

4.9 12.1

S O H + Ag + ~ [ S O - - A g +] + H + S O H + Ag + m [ S O - - A g O H °] + 2H +

Journal of Colloid and Interface Science, Vol. 67, No. 1, October 15, 1978

104

DAVIS AND LECKIE

Table III gives a compilation of intrinsic surface complexation constants for several metal ions and hydrous oxides. These constants were obtained by fitting model calculations with experimental adsorption data on oxides with known surface parameters (Table I). Note the greater stability of bonds formed by hydroxo complexes in all cases as compared to free aquo metal ions. With the present state of knowledge concerning the nature of bonds formed at the oxide/ water interface, it is not possible to attribute this difference in stability to either physical or specific chemical energy terms. A plausible explanation is that metal ions are more easily hydrolyzed at the surface than in bulk water due to changes in the entropic environment (37). Since the only physical parameter we have defined explicitly is a coulombic energy term, the stability constants may include the effect of other physical phenomena as well as specific chemical bonding energy. For example, these results are in agreement with the model of James and Healy (10), who stated that adsorption of lower-charged hydrolyzed metal ions was enhanced over that of aquo metal ions because of differences in solvation energy. Among the cations shown, Pb(II) and Zn(II) form stronger bonds with oxide surfaces than Cu(II), Cd(II), and Ag(I), but the differences are not large. However, this does not imply that the adsorption edges of Pb(II) or Zn(II) for a given oxide substrate will occur at a lower pH, since adsorption density is a complex function of solution and surface speciation. The surface complexation constants for each metal ion are of the same order of magnitude for most oxides, e.g., Cd(II) with TiO2, a-FeOOH, y-AI2Oz, and Fe~Oa" H20(am), with the exception of or-quartz. Similar to the results found for major electrolyte cations (1), we find that the intrinsic constants (K~ t) of heavy metal ions with a silica surface are considerably smaller than with other oxide surfaces. Thus, it would seem more appropriate to use intrinsic constants for an oxide other than Si02 as estimates of surface sta-

bility constants needed for predictive calculations of trace metal distribution in most natural aqueous systems. Benjamin (35) has found that the pH adsorption edges of several metal ions, e.g., Pb(II), Cd(II), on some hydrous oxides (3/A12Oa, Fe~O~'H~O(am)) are dependent on the total metal concentrations, even at a low adsorption density. This phenomenon was attributed to the existence of an array of oxide surface sites of variable complexation strength for metal ions. Thus, to achieve agreement between Benjamin's experimental data and model calculations, the surface complexation constants must vary as a function of the total metal concentration; hence, we have included this experimental parameter in Table III. While we have not studied this phenomenon in detail, it is apparent that the value of log /~PbOH +r~int decreases monotonically with increasing total Pb(II) concentration (Table III). The implication for predictive calculations of the equilibrium distribution of trace metals in natural aquatic systems is that the best estimates of surface stability constants are obtained using experimental data for systems with very dilute total metal concentrations (10 -8 to 10-6 M). Most experimental adsorption data have been reported for higher total metal concentrations (5-7); however, more recent investigations (15, 35, 36) have included systems with total metal concentrations typical of natural waters.

Charge Distribution and Coordination Chemistry at the Oxide~Water Interface The model calculations presented involve various assumptions necessary to write surface equilibria for adsorbing metal ions. For example, Eq. [10] implies that we may simulate the coulombic energy contribution to formation of the surface complex, S O - PbOH +, by the term exp[(e$0 - etka)/kT]. Thus, we have assumed that the charge distribution with respect to an adsorbed monohydroxometal(II) ion can be modeled by charges of - 1 and +1 in the tr 0 and tro planes, respectively. While this assumption

Journal of Colloid and Interface Science, Vol. 67, No. 1, October 15, 1978

ADSORPTION

105

OF METAL IONS

TABLE III Intrinsic Surface Complexation Constants for Metal Ions and Oxide Substrates Metal: Cd(H)

Surface reactions:

SO- + Cd2+ s=~SO-_Cd z+ SO- + CdOH+ ~ SO--CdOH +

log K ~ +

log KlY~oa+

Total metal concentration (M)

TiO2

7.2

10.1

2 x 10 -4

a-FeOOH

6.0

9.3

2 × 10 -4

(7)

y-Al2Oa FezO3'H20(am) a-Quartz

5.9 5.8 a

9.7 9.8 5.0

5 x 10 -~ 5 x 10 -7 5 X 10 - r

(35) (35) (35)

Oxide

Metal: Pb(ll)

Oxide

log K ~ +

Surface reactions:

log Kgtoa+

Total metal concentration (M)

K tat estimated

from Ref. (7) Fig. 7 o f this s t u d y

SO- + 1 ~ + ~; S O - - I ~ + SO- + PbOH+ s:; SO-_PbOH +

K~t estimated from Ref.

FezO3"H20(am)

6.9

11.1

5 × 10 -7

(35)

y-A1203

6.4

10.2

5 x 10 -7

(35) (35)

y-A1203

'~

8.7

5 x 10 -5

T-AI203

6.4

8.9

2.9 x 1 0 - '

a-Quartz

4.6

7.8

5 x 10 -7

Metal: Cu(lI)

Surface reactions:

Total metal concentration

(5) Fig. 6 o f this s t u d y (36) SO- + Cu*+ s=~SO-_Cu2+ SO- + CuOH+ s:~ SO-_CuOH +

/ ~ t estimated from Ref.

Oxide

int log Kc~a+

tat log Keaoa+

(M)

T-A1203 Fe2Oa" H ~ O ( a m )

6.4 6.5

9.7 9.6

10 -e 10 -e

(35) (15) Fig. 8 o f this s t u d y

a-Quartz

2.8

6.6

10 -6

(36)

Metal: Zn(H)

Oxide y_Al203 Fe203"H20(am) a-Quartz

Surface reactions:

log K~2+

log K~2~on+

Total metal concentration (M)

a a a

10.1 9.4 7.5

5 X 10 -7 5 × 10 -7 I0 -s

Metal: Ag(1)

Surface reactions:

Oxide

log K ~t+

log K~oa

Total metal concentration (M)

Fe203"H~O(am)

5.7

10.4

4 x 10 -7

c¢-Quartz

3.2

6.1

4 x 10 -~

SO- + Zn~+ ~ SO--Zn 2+ SO- + ZnOH+ s=~SO-_ZnOH+

Kt*t estimated from Ref. (35) (35) (43) SO- + Ag + ~ SO--Ag+ SO- + AgOH s:; SO--AgOH

Klm estimated from Ref. (15) Fig. 10 o f this s t u d y (15, 16)

a U n a b l e t o d e t e r m i n e s i n c e m o d e l c a l c u l a t i o n s w e r e a d e q u a t e w i t h t h e m o n o h y d r o x o c o m p l e x as t h e o n l y adsorbing species. Journal o f CoUoid and Interface Science, Vol. 67, No. I, October 15, 1978

106

DAVIS AND LECKIE

contributes significantly to the mathematical simplicffy of our site-binding model, its validity and implications for surface bonding of adsorbed metal ions need to be discussed. The exponential term in Eq. [10] suggests that the proton released from the hydration sphere of pb2+(aq) originates from a cis position relative to coordination by the oxide surface site, resulting in a charge of +1 in the tra plane. For metal ions which undergo ligand substitution slowly, e.g., Co(III) in the ammine system, the influence of a ligand in affecting the acidity o f a trans or cis water molecule can be assessed. This is much more difficult experimentally for labile metal ions, e.g., Cu(II) or Pb(II), and the effects observed for robust complexes are not well enough understood to provide a basis for predicting the results (H. Taube, personal communication). Thus, it is equally possible that the expelled proton originates from a trans position and that the coulombic energy term in Eq. [10] should be written in terms of a charge of +2 in the o-8 plane and - 1 in the trd plane, i.e., exp[(e~b0 - 2e~b~ + e~a)/kT]. However, the net change of this modification to Eq. [10] is [(e~bd --ed/~)/kT], and the overall effect on calculated adsorption density of metal ions in dilute solution is relatively minor. This occurs because the variation in I* 1 and I,dl with pH in 0.1 M electrolyte is relatively small compared to that of [00[ (1). As a consequence, we cannot predict at this time from which coordination position the proton will be expelled. Nonetheless, our conclusions regarding surface speciation of adsorbed metal ions remain unchanged, i.e., adsorption of hydrolytic complexes is more consistent with experimental adsorption data than complexation by bidentate surface sites. We have examined elsewhere (37) the effects of other modifications to the charge distribution within the interfacial region. For example, it is possible to obtain agreement between calculated and experimental adsorption density using equilibria that consider complexation of metal ions by bidenJournal of CoUoid and Interface Science, Vol. 67, No. 1, October 15, 1978

tate surface sites, if one assumes the existence of partial (fractional) charges in the o-0 and o-s planes. This may be a reasonable hypothesis since the occurrence of nephelauxetic effects in coordination compounds indicate that the actual charges of coordinated metal ions are smaller than the formal oxidation number (44). In an even more general approach, one might consider integral or fractional charges located in an intermediate plane between the or0 and o-s planes. Since the appropriate corrections are difficult to estimate, however, we prefer the simpler approach of integral charges located in the tr0 and tr~ planes, as presented in this paper. Refinement of the model awaits the results of further experimental work regarding both the physical and the chemical nature of surface bonding of metal ions at the interface. An important advantage of this adsorption model is that surface charge and diffuse layer potential are also calculated and may be compared with potentiometric titration results or electrokinetic measurements, if available. Thus, the effects of adsorbing ions on the interfacial electrostatic field and colloid stability can be predicted. While simpler, strictly chemical models may be more practical for some applications, e.g., modeling trace metal adsorption in marine waters, we have shown that a more complete physicochemical description of the oxide/water interface is useful for determining the stoichiometry of surface reactions and the speciation of adsorbed metal ions. ACKNOWLEDGMENTS This work was partially supported by Air Force Contract F29601-75-C-0028 and EPRI Contract RP-910. James A. Davis received partial support from Department of Civil Engineering Fellowship funds. The authors wish to thank P. W. Schindler, H. Taube, R. O. James, W. Stumm, and G. A. Parks for discussion and criticism. We would also like to express our appreciation to J. Westall and F. M. M. Morel for assistance with the computer program, MINEQL. REFERENCES 1. Davis, J. A., James, R. O., and Leckie, J. O., J. Colloid Interface Sci. 63, 480 (1978).

ADSORPTION OF METAL IONS 2. Westall, J. C., Zachary, J. L., and Morel, F. M. M., Technical Note No. 18. Water Quality Laboratory, Department of Civil Engineering, Massachusetts Institute of Technology, Cambridge, Mass., 1976. 3. Davis, J. A., and Leckie, J. O., submitted for publication. 4. Loganathan, P., Burau, R. G., and Fuerstenau, M. C., Proc. Soil Sci. Soc. Amer. 41, 57 (1977). 5. Hohl, H., and Stumm, W. ,J. Colloidlnterface Sci. 55, 281 (1976). 6. Schindler, P. W., Fiirst, B., et al., J. Colloid Interface Sci. 55, 469 (1976). 7. James, R. O., Stiglich, P. J., and Healy, T. W., Faraday Discuss. Chem. Soc. 59, 142 (1975). 8. MacNaughton, M. G., and James, R. O.,J. Colloid Interface Sci. 47, 431 (1974). 9. Murray, J. W., Geochim. Cosmochim. Acta 39, 505 (1975). 10. James, R. O., and Healy, T. W., J. Colloid Interface Sci. 40, 65 (1972). 11. Bowden, J. W., Bolland, M. D. A., et al., Nature Phys. Sci. 245, 81 (1973). 12. Stumm, W., Hohl, H., and Dalang, F., Croat. Chem. A c t s 48, 491 (1976). 13. Dugger, D. L., Stanton, J. H., et al., J. Phys. Chem. 68, 757 (1964). 14. Jenne, E. A., Advan. Chem. Set. 73, 337 (1968). 15. Davis, J. A., "Adsorption of Trace Metals and Complexing Ligands at the Oxide/Water Interface," Ph.D. Thesis. Stanford University, Stanford, Calif., 1977. 16. Davis, J. A., and Leckie, J. O., submitted for publication. 17. Murphy, P. J., Posner, A. M., and Quirk, J. P., J. Colloid Interface Sci. 56, 270 (1976). 18. Dousma, J., and de Bruyn, P. L., J. Colloid Interface Sci. 56, 527 (1976). 19. Van der Giessen, A. A., J. Inorg. Nucl. Chem. 28, 2155 (1966). 20. Yates, D. E., "The Structure of the Oxide/Aqueous Electrolyte Interface," Ph.D. Thesis. University of Melbourne, Melbourne, Australia, 1975. 21. Gast, R. G., Landa, E. R., and Meyer, G. W., Clays Clay Minerals 22, 31 0974). 22. Avotins, P. V., "Adsorption and Coprecipitation Studies of Mercury on Hydrous Iron Oxides," Ph.D. Thesis. Stanford University, Stanford, Calif., 1975. 23. Van den Hul, and Lyklema, J., J. Amer. Chem. Soc. 90, 3010 (1968). 24. Parks, G. A., and de Bruyn, P. L., J. Phys. Chem. 66, 967 (1962). 25. Yates, D. E., and Healy, T. W., J. Colloid Interface Sci. 52, 222 (1975).

107

26. Parks, G. A., Chem. Rev. 65, 177 (1965). 27. James, R. O., Davis, J. A., and Leckie, J. O., J. Colloid Interface Sci., 65, 331 (1978). 28. Iler, R. K., in "Surface and Colloid Science" (E. Matijevi~, Ed.), Vol. 6. Wiley-Interscience, New York, 1973. 29. Yates, D. E., Levine, S., and Healy, T. W., Chem. Soc. Faraday Trans. 1 70, 1807 (1974). 30. Stumm, W., Huang, C. P., and Jenkins, S. R., Croat. Chem. Acts 42, 223 (1970). 31. Jenne, E. A., in "Symposium on Molybdenum in the Environment" (W. Chappel and K. Petersen, Eds.), Vol. 2, Chap. 5. Marcel Dekker, New York, 1977. 32. Schindler, P. W., and Gamsj~iger, H., KolloidZ. Z. Polym. 250, 759 (1972). 33. Schindler, P. W., and Kamber, H. R., Helv. Chim. Acta 51, 1781 (1968). 34. Stiglich, P. J., "Adsorption of Cadmium(II) Complexes at the Oxide/Water Interface," M.S. Thesis. University of Melbourne, Melbourne, Australia, 1976. 35. Benjamin, M. M., "Adsorption of Dissolved Cadmium, Zinc, Copper, and Lead on Oxide Surfaces in Model Natural Systems," Ph.D. Thesis. Stanford University, Stanford, Calif., 1978. 36. Vuceta, J., "Adsorption of Pb(II) and Cu(II) on a-Quartz from Aqueous Solutions: Influence of pH, Ionic Strength, and Complexing Ligands," Ph.D. Thesis. Calif. Institute of Technology, Pasadena, Calif., 1976. 37. Davis, J. A., and Leckie, J. O., in "Chemical Modeling-Speciation, Sorption, Solubility, and Kinetics in Aqueous Systems" (E. A. Jenn6, Ed.) ACS Symposium Series, in press. 38. Huang, C. P., "The Chemistry of the Aluminum Oxide-Electrolyte Interface," Ph.D. Thesis. Harvard University, Cambridge, Mass., 1971. 39. Per'i, J. B., J. Phys. Chem. 69, 211 (1965). 40. Hingston, F. J., Posner, A. M., and Quirk, J. P., Advan. Chem. Set. 79, 82 (1968). 41. Abendroth, R. P.,J. Colloid Interface Sci. 34, 591 (1970). 42. Armistead, C. G., Tyler, A. J., et al., J. Phys. Chem. 73, 3947 (1969). 43. James, R. O., and MacNaughton, M. G., Geochim. Cosmochim. Aeta 41, 1549 (1977). 44. J0rgensen, C. K., "Absorption Spectra and Chemical Bonding in Complexes." AddisonWesley, Reading, Mass., 1962. 45. Smit, W., Holten, C. L. M., et al., J. Colloid Interface Sci., 63, 120 (1978). 46. Smit, W., and Stein, H. N.,J. Electroanal. Chem., 91, 393 (1978).

Journal of CoUoidand InterfaceScience. Vol.67, No. 1, October15, 1978