Aquatic Chemistry - American Chemical Society

1 The Inner-Sphere Surface Complex Downloaded by UCSF LIB CKM RSCS MGMT on November 25, 2014 | http://pubs.acs.org Publication Date: May 5, 1995 | doi: 10.1021/ba-1995-0244.ch001

A Key to Understanding Surface Reactivity Werner Stumm Swiss Federal Institute of Technology, Zürich; EAWAG (Institute for Environmental Science and Technology), CH-8600, Dübendorf, Switzerland

Functional groups on the interface of natural solids (minerals and particles) with water provide a diversity of interactions through the formation of coordinate bonds with H , metal ions, and ligands. The con+

cept of active surface sites is essential in understanding the mechanism of many surface-controlled processes (nucleation and crystal growth, biomineralization, dissolution and weathering of minerals, soil formation, catalysis of redox processes, and photochemical reactions). The enhancement of the dissolution rate by a ligand implies that surface complex formation facilitates the release of ions from the surface to the adjacent solution. These ligands bring electron density within the coordinating sphere of the central ion. Surface species thus destabilize the bonds in the surface lattice; they are especially efficient in the dissolution of iron and aluminum oxides and of aluminum silicates. Ascorbate, phenols, and S(-II) compounds, including H S, readily form 2

surface complexes with Fe(III) or Mn(III,IV)

(hydr)oxides that sub-

sequently undergo electron transfer and the release of Fe(II) or Mn(II) into solution. Reductive and nonreductive dissolutions are markedly inhibited by competitive (ligand exchange) adsorption of inorganic oxoanions. These oxoanions can form bi- or multinuclear surface complexes. A better understanding of the electronic structure of the interface of solids and aquatic solutes would push the boundaries of aquatic surface chemistry.

I N T E R A C T I O N AT T H E S O L I D - W A T E R I N T E R F A C E can be characterized i n terms o f the c h e m i c a l a n d physical properties o f water, the solute, a n d the

0065-2393/95/0244-0001$09.28/0 © 1995 American Chemical Society

In Aquatic Chemistry; Huang, C., et al.; Advances in Chemistry; American Chemical Society: Washington, DC, 1995.

Downloaded by UCSF LIB CKM RSCS MGMT on November 25, 2014 | http://pubs.acs.org Publication Date: May 5, 1995 | doi: 10.1021/ba-1995-0244.ch001

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AQUATIC CHEMISTRY

sorbent. The two basic processes in the reaction of solutes with natural sur faces are the formation of coordinate bonds (surface complexation) and hy drophobic adsorption. Hydrophobic adsorption is primarily driven by the incompatibility of nonpolar, hydrophobic substances with water. The formation of coordinate bonds is based on the generalization that the solids can be considered either inor ganic or organic polymers; their surfaces can be seen as extending structures bearing surface functional groups. These functional groups contain the same donor atoms found in functional groups of solute ligands such as - O H , - S H , - S S , and - C 0 H . Such functional groups provide a diversity of interactions through the formation of coordinate bonds. Similarly, ligands can replace sur face O H groups (ligand exchange) to form ligand surface complexes. The concept of active sites has helped explain catalysis by enzymes and coenzymes. Although surface functional groups are less specific than enzymes, they form an array of surface complexes whose reactivities determine the mechanism of many surface-controlled processes. Many mechanisms can be described readily in terms of Br0nsted acid sites or Lewis acid sites. O f course, the properties of the surfaces are influenced by the properties and conditions of the bulk structure, and the action of special surface structural entities will be influenced by the properties of both surface and bulk. List I gives an overview of the major concepts and important applications. Surface chemistry of the oxide-water interface is emphasized here, not only because the oxides are of great importance at the mineral-water (includ ing the clay-water) interface but also because its coordination chemistry is much better understood than that of other surfaces. Experimental studies on the surface interactions of carbonates, sulfides, disulfides, phosphates, and biological materials are only now emerging. The concepts of surface coordi nation chemistry can also be applied to these interfaces. This chapter is designed 2

• to briefly review surface complexation theory, reflecting on the nature of site-specific binding to H , metal ions, and ligands +

• to discuss the need for assessing the bonding between solids and solutes to understand better the reactivity of the solid-water i n terface and to illustrate this reactivity in terms of surface-con trolled dissolution of oxides and silicates • to present exemplifying experimental evidence on various factors that enhance or inhibit dissolution to make the point that we need a better appreciation of the electronic structure and the geometry of the bonding at the solid-water interface to predict reactivity • to exemplify some applications of the effects of surface complex formation, surface reactivity enhancement, and inhibition of dis-


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The Inner-Sphere Surface Complex

List I. Coordination Chemistry of the Solid-Water Interface: Concepts and Applications in Natural and Technical Systems

Surface Complex Formation Interaction with H , OH" Metal ions Ligands (ligand exchange)


+

Thermodynamics of Surface Complex Formation Κ (mass law constants, corrected for electrostatic effects AG, ΔΗ Kinetics of Surface Complex Formation Rates of sorption and desorption Structure of Surface Compounds (Surface Speciation) Inner-sphere versus outer-sphere Mononuclear versus binuclear Monodentate versus bidentate Establishment of Surface Charge (Structure of Lattice) Defect sites Adatoms, kinks, steps, ledges Lattice statistics Microtopography

Applications: Distribution of Solutes between Water and Solid Surface Binding of Reactive Elements to Aquatic Particles in Natural Systems Regulation of metals in soil, sediment, and water systems Regulation of oxyanions of P, As, Se, and Si in water and soil systems Interaction with phenols, carboxylates, and humic acids Transport of reactive elements including radionuclides in soils and aquifers Binding of Cations, Anions, and Weak Acids to Particles in Technical Systems Corrosion, passive films Processing of ores, flotation Coagulation, flocculation, filtration Ceramics, cements Photoelectroehemistry (electrodes, oxide electrodes, and semiconductors) Surface Charge Resulting from the Sorption of Solutes Particle-particle interaction; coagulation, filtration

Applications: Rate Dependence on Surface Speciation Natural Systems Dissolution of Oxides, Silicates, Carbonates, and Other Minerals Weathering of minerals Proton- and ligandpromoted dissolution Reductive dissolution of Fe(III) and Mn(III,IV) oxides Formation of Solid Phases Heterogeneous nucleation Surface precipitation, crystal growth Biomineralization Surface-Catalyzed Proceses ( Photo)redox processes Hydrolysis of esters Transformations of organic matter by F e and M n (photo)redox cycles Oxygenation of Fe(ll), Mn(ll), Cu(I), and V(IV) Technical Systems Passive films (corrosion) Photoredox processes with colloidal semiconductor particles as photocatalyst (e.g., degradation of refractory organic substances) Photoelectroehemistry (e.g., photoredox processes at semiconductor electrodes)

SOURCE: Modified from reference 1.


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AQUATIC CHEMISTRY

solution in natural weathering processes, in heterogeneous pho tochemical processes, and in technical systems (corrosion and dissolution of passive iron oxide films)

Surface Coordination Chemistry


Inner- and Outer-Sphere Complexes.

As illustrated in Figure 1,

a cation can associate with a surface as an inner-sphere or an outer-sphere complex, depending on whether a chemical bond is formed (i.e., a largely covalent bond between the metal and the electron-donating oxygen ions, as in an inner-sphere complex) or whether a cation of opposite charge approaches the surface groups within a critical distance. As with solute ion pairs, the cation and the base are separated by one or more water molecules (J, 2). Further more, ions may exist in the diffuse swarm of the double layer.

Figure 1. Fart a: Surface complex formation of an ion (e.g., cation) on a hydrous oxide surface. The ion may form an inner-sphere complex (chemical bond), an outer-sphere complex (ion pair), or be in the diffuse swarm of the electnc double layer. (Reproduced with permission from reference 2. Copyright 1984.) Part b: Schematic portrayal of the hydrous oxide surface, showing planes associated with surface hydroxyl groups (s), inner-sphere complexes (a), outer-sphere complexes (β), and the diffuse ion swarm (d). In the case of an inner-sphere complex with a ligand (e.g., F~ or ΉΡΟ/'), the surface hydroxyl groups are replaced by the ligand (ligand exchange). (Modified from reference 3.)



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It is important to distinguish between outer-sphere and inner-sphere com plexes. In inner-sphere complexes the surface hydroxyl groups act as σ-donor ligands, which increase the electron density of the coordinated metal ion. Cu(II) bound in an inner-sphere complex is a different chemical entity from Cu(II) bound in an outer-sphere complex or present in the diffuse part of the double layer. The inner-spheric Cu(II) has different chemical properties; for example, it has a different redox potential with respect to Cu(I), and its equa torial water is expected to exchange faster than that in Cu(II) bound in an outer-sphere complex. As we shall see, the reactivity of a surface is affected, above all, by inner-sphere complexes. List II summarizes schematically the type of surface complex formation equilibria that characterize the adsorption of H , OH", cations, and ligands at a hydrous oxide surface. The various surface hydroxyls formed at a hydrous oxide surface may not be fully equivalent structurally and chemically. However, to facilitate the schematic representation of reactions and of equilibria, we will consider the chemical reaction of surface hydroxyl group, S - O H . The following surface groups can be envisaged. +

OH , S - O H

.OH S—OH OH

SCT n u

These functional groups have donor properties similar to those of their counterparts in dissolved solutes such as hydroxides or carboxylates. Thus, List II. Adsorption (Surface Complex Formation Equilibria) Acid-base equilibria S-OH S-OH

+ H

(+

+

* = 5 S-OH

2

+

H 0)

O H " ) ± = * S - O " (+

2

Metal binding + M

± = 5

S-OM ^"

M

* = 5

(S-0) M ^

H 0 2

±=*

S-OMOH ~"

+ L-

* = 5

S-L +

t=*

S -L

S-OH

z+

2 S-OH + S-OH

+

M

3 +

z+

+

(

(

2

+

1 ) +

H

2H

+

2 ) +

(

+

2 ) +

+

+ 2 H

+

Ligand exchange (L~ = ligand) S-OH

2 S-OH +

L"

2

+

O H " + 2 O H "

Ternary surface complex formation S-OH S-OH

+ L - + M~' +

L" +

M

=

+

S-L-M

+

t=5

5 +

S-OM-L

+ ( 3

"

O H " 2 ) +

+

H

+

SOURCE: Modified from reference 4.


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AQUATIC CHEMISTRY

deprotonated surface groups (S-O") behave like Lewis bases and the sorption of metal ions (and protons) can be understood as competitive complex formation.

The adsorption of lig

Adsorption of Ligands on Metal Oxides.

ands (anions and weak acids) on metal oxide and silicate surfaces can also be compared with complex formation reactions in solution.


Fe(OH)

2+

+ F"

S-OH + F -

• FeF

2 +

+ OH"

(la)

• S-F + O H "

(lb)

The central ion of a mineral surface acts as a Lewis acid and exchanges its structural O H with other ligands (ligand exchange). In this case consider the surface of Fe(III) oxide as an example. S - O H corresponds to = F e - O H . A Lewis acid site is a surface site capable of receiving a pair of electrons from the adsorbate. (A Lewis base site has a free pair of electrons—like the oxygen donor atom in a surface O H " group—that can be transferred to the adsorbate.) The extent of surface complex formation (adsorption) for metal ions and an ions is strongly dependent on p H and on the release of protons and O H " ions, respectively. In addition to monodentate surface complexes, bidentate (mon onuclear or binuclear) surface complexes can be formed. 2S-OH + Cu

-S-OH I -S-OH

9

+ Cu * 2

2 +

> (S-0) Cu + 2 H 2

-S-Ov I Cu + 2 H -S-c/ x

o -

c ' ° "

sFe'

=FeOH I H PO; sFeOH

=Fe — C \ , O" I Ρ ,Fe-0/ *0

+

2

N

0 —

(2b)

+

sFeOH + H C 0 ; (Oxalate) 2

(2a)

+

I

+H 0 2

( 3 )

^

(4)

The following criteria characterize all surface complexation models (5): Sorption takes place at specific surface coordination sites. Sorption reactions can be described by mass law equations.


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• Surface charge results from the sorption (surface complex for mation) reaction itself.


• The effect of surface charge on sorption (the extent of complex formation) can be taken into account by applying to the mass law constants for surface reactions a correction factor derived from the electric double-layer theory. The extent of adsorption, or surface coordination, and its p H dependence can be accounted fpr by mass law equilibria (Figure 2). Their equilibrium constants reflect the affinity of the surface sites for H , metal ions, and ligands. The tendency to form surface complexes may be compared with the tendency to form corresponding (inner-sphere) solute complexes (4-6). Figure 3 shows the relation between the solute complex formation of F e O H or A l O H with various ligands and the surface complexation of = F e O H and = A l O H surface groups with the same ligands. The reasonably good correlation obtained in this and similar linear free energy relation ( L F E R ) plots (4-6) indicates that the same chemical mode of interaction occurs in solution and at the surface and that the available sorption data are consistent with one another. Therefore, such L F E R s may be used to predict intrinsic sorption constants from solute complex formation constants and vice versa. +

2

Surface Complex Formation on Carbonates.

+

2 +

There are various

possibilities for functional groups on the surface of carbonates, sulfides, phos phates, and similar compounds. By using a very simple approach similar to the one used for hydrous oxides (chemisorption of H 0 ) , one could postulate surface groups for carbonates (e.g., F e C 0 ) as shown in List III. As indicated in Scheme I, it is reasonable to assume that H , O H " , H C 0 ~ , C0 (aq), and F e can interact with MC0 (s) and affect its surface charge. Surface complex formation of the surface groups with ligands and metal ions can occur (9). 2

3

+

2

2 +

3

3

Surface Reactivity Dependence on Surface Structure Many heterogeneous processes such as dissolution of minerals, formation of the solid phase (precipitation, nucleation, crystal growth, and biomineralization), redox processes at the solid-water interface (including light-induced reactions), and reductive and oxidative dissolutions are rate-controlled at the surface (and not by transport) (10). Because surfaces can adsorb oxidants and reductants and modify redox intensity, the solid-solution interface can catalyze many redox reactions. Surfaces can accelerate many organic reactions such as ester hydrolysis (11). The mechanisms of most surface-controlled processes depend on the co ordination environment at the solid-water interface. Above all, they depend



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Figure 2. These curves were calculated with the help of experimentally deter mined equilibrium constants. Part a: Extent of surface complex formation as a function of pH (measured as mole percent of the metal ions in the system, ad sorbed or surface-bound). Total ion concentration [TOTFe] = lOr* M (2 Χ ΙΟ' mol I h of reactive sites; metal concentrations in solution = 5 Χ 10~ M; I = 0.1 Μ NaN0 . (The curves are based on data compiled by Dzombak and Morel in reference 5.) Part b: Surface complex formation with ligands (anions) as a func tion of pH. Binding of anions from dilute solutions (5 Χ 10~ M) to hydrous ferric oxide; [TOTFe] = lOr* M. I = 0.1. (Curves are based on data from Dzom bak and Morel in reference 5.) Part c: Binding of phosphate, silicate, and fluoride on goethite (a-FeOOH); the species shown are surface species (6 g/Lof FeOOH, P = 10 M, Si = 8 X 10~ M). (Reproduced with permission from reference 4

7

3

7

T

3

T

4

6. Copyright 1981.)


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20

1

1

1

9

1

1

1

20

Λ

I

1

Aquatic Chemistry - American Chemical Society

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