C0LLOIDS A AND Colloids and Surfaces ~URFACE~

ELSEVIER A: Physicochemical and Engineering Aspects 120 (19971143 166

Reactivity at the mineral-water interface: dissolution and inhibition

Werner Stumm

lnstitute for Environmental Science and Technology ( EAWAG), Swiss Federal Institute ~?]Techm)logy ( I£TH ZtTrich), CH-8600 D~Tbendorjl Switzerland

Received 2 October 1995: accepted 10 June 1996

Abstract

The reactivity of the mineral-water interface is often interpreted with the help of models used in electrochemistry, in solution coordination chemistry and in crystallography. Progress in understanding mechanisms of growth and dissolution of crystals and the inhibition of these processes depends on a better integration of these models. It is shown that dissolution can be explained in terms of a ligand exchange process: simplified rate laws for proton- and ligand-protonated dissolution rates, being related to surface bound protons and ligands, respectively, can be derived. Further refinement in interpreting surface reactivity comes from an appreciation of the molecular structures at the mineral water interface; here significant advances have been made by X-ray absorption spectroscopy, especially EXAFS (Extended X-ray Absorption Fine Structure spectroscopy), which permit distinction to be made between outer-sphere and inner-sphere surface complexes, and in many cases to determine the structure of the surface species at different crystallographic planes (e.g., bi-nuclear or mono-nuclear linkage of ligands on metal ions to surface metal centers). Such information coupled with solution-chemical studies on the extent of adsorption can provide new insight into the mechanisms of dissolution reactions and their inhibition and surface poisoning. A few experimental results are given to exemplify the factors that enhance and inhibit the non-reductive (EDTA) and reductive dissolution (by HzS) of Fe(IlI)(hydr)oxides. Binuclear surface complexes by multivalent cations and by oxoanions, such as phosphate, arsenate and borate, are believed to be efficient inhibitors for oxide dissolution because they form bi- or multinuclear inner- sphere surface complexes that can bridge two or more metal centers in the surface lattice; the simultaneous removal of such bi- or multinuclear surface complexes from the surface is energetically unfavorable. Proton and ligand promoted dissolution reactions and their inhibition by oxoanions and bi- or multinuclear surface complexes are not only relevant in geochemistry (weathering, soil-formation, transfer of elements and pollutants) but also in metallic corrosion, formation and breakdown of passive oxide films.

Keywords: Corrosion inhibition; Dissolution; Ligand exchange; Mineral: Oxide: Passivation

1. Introduction

Most geochemical processes involve the forma- t ion (prec ip i ta t ion) and the d isso lu t ion (weather- ing) of the solid phase. These processes are i m p o r t a n t in soil so lu t ion chemis t ry and in the geochemis t ry of na tu ra l waters. Fig. 1 gives a

survey of the d isso lu t ion rates of a few selected minerals . It i l lustrates the var ia t ion in d isso lu t ion rates of different minerals by many orders of magn i tude and shows a p ronounc e d pH depen- dence; typical ly, the d isso lu t ion rates of oxides and silicates below their p H of zero poin t of charge increase with decreas ing pH and increase with

0927-7757/97 ~$17.00 Copyright ~) 1997 Elsevier Science B.V. All rights reserved PIL- S0927-7757(96)03866-6

144 W. Stumm/Colloids Surfaces A: Physicochem. Eng. Aspects 120 (1997) 143-166

10-5

10-7

"6 g

• ~ 10_9

10-11

10-13

4 6 8 10 pH

xln [h) 0.001

0.01

0.1

1

10

102

103

104

105

Fig. 1. Survey on dissolution rates of minerals. For experimen- tal data see: Calcite and Dolomite: Chou and Wollast [1]; Anorthite (CaAIzSi208): Amrhein and Suarez [2]; Forsterite (MgzSiO4): Blum and Lasaga [3] Biotite (micagroup): Lin and Clemency [4] (more recent data by Acker and Bricker, 1992: their rates that are ca. 1 order of magnitude smaller); ct-FeOOH (Goethite): Zinder et al. [5]; ~-A1203: Furrer and Stumm [6]; Kaolinite, Muscovite: Stumm and Wieland I-7 ]; Quartz: Wollast and Chou I-8]; Albite (NaAISi3Os): Chou and Wollast [9]. On the right hand scale is the half life (hours) given for the functional surface groups -=S OH, S CO3 H. (The experimental dissolution rate, Ru, (mol area -1 t ime- l ) is proportional to the total concentration of surface sites, S t, (mol area- a), RH = kmS~, where k,, (time -a) represents the mean reactivity of all surface sites. The half life of a surface site is given by ~a/2 = In 2~kin. A surface density of 10 surface sites per nm 2 have been assumed.)

increasing pH in the alkaline region. Several path- ways contribute to the overall dissolution reac- tions. The understanding of the dynamics of the dissolution processes and their retardation (inhibi- tion) requires an appreciation of the reactions which describe the transfer of the chemical species between the mineral and the aqueous solution and an understanding of the structure and chemical bonding at the mineral water interface.

1.2. Objectives

In this discourse, we review the various factors that influence the dissolution of minerals and their inhibition. Because we lack sufficient knowledge on the ways molecules and ions interact at the mineral-water interface, above all, on the electronic structure of the bonding between solids and solutes, we will illustrate how mechanisms of the reactivity of the hydrous mineral surface can be derived from models typically encountered, (1) in electrochemis- try, (2) in solution coordination chemistry and (3) in crystallography.

Such field-specialized models have been used in the literature to account for mechanisms of growth and dissolution of crystals and the inhibition of these processes, but one of the objectives of this paper is to illustrate that a more unifying under- standing could be obtained, if we could translate the specialized treatments into a more comprehen- sive general picture.

Progress has been made in the last one to two decades above all, (1) in the understanding of surface complex formation as a treatment of the interaction of reactive species between the solution and the mineral surface [ 10-14] and (2) in techno- logical advances of in situ spectroscopic techniques (especially, high flux synchroton radiation and high counting rate fluorescence X-ray detectors) to studying molecular structures at the particle water interface [ 15-18]. The structure of various surface complexes upon adsorption has been described for various laboratory model systems as well as for natural particles; such descriptions provide a struc- tural chemical appreciation of the simple assump- tions made by thermodynamic models and the information which can be deduced at the atomic level, can, when coupled with the solution chemical studies, provide new insight into the mechanisms of dissolution reactions and their inhibition and crystal growth and particle growth poisoning.

This is not a comprehensive literature review; instead a few selected case examples are used to illustrate pertinent mechanisms. Emphasis is given on the interpretation of mineral dissolution as a ligand exchange reaction and on the structural surface chemistry of hydrous oxides. Although the ideas presented are believed to be valid for all

g4 Stumm/Colloids Surfaces A: Ph.vsicochem. Eng. Aspects 120 [ 1997) 143 166 145

types of minerals, including carbonates [19,20], sulfides [21] and phosphates [22], we restrict our discussion to silicates and oxides. Among the latter minerals, Fe(III)(hydr)oxides are often used for exemplification because most information is avail- able for these compounds.

2. Molecular models for the mineral-water interface

Three models have been instrumental in bringing about an appreciation of the interaction and reac- tivity at the mineral-water interface: ( 1 ) the mineral surface as an electrode: dissolution as an electro- chemical process and an electric double layer to account for the charge separation at the interface; (2) coordination chemistry at the mineral surface and at the solid-water interface; (3) crystal chemis- try at the surface, molecular structure and surface morphology.

2.1. Electric double layer

At the mineral water interface there exists a separation of electric charge. The distribution of ions (charges) in the neighborhood of a charged surface is idealized as an electrochemical double layer; one layer is envisaged as a fixed charge or surface charge attached to the solid surface while the other layer is distributed more or less uniformly on the electrolyte or contact with the charged surface. The dissolution of an oxide has been interpreted as an electrochemical process [23,24]. The rates of adsorption and desorption of reactive species and the rates of dissolution and precipita- tion depend on the surface charge of the mineral surface.

2.2. Sut~[hce coordination model

While the surface in the electric double layer model is assumed to be a structure-less continuum which interacts with the solution only by its electric charge, the basic concept in the surface coordina- tion model are the surface functional groups formed on all inorganic solids; the surface func- tional group of greatest abundance on hydrous

minerals is the hydroxy group (see Fig. 2a). Specific adsorption occurs through coordinative inter- actions. The surface moiety of a mineral surface may be looked at as a metal complex with aquo, oxygen, hydroxo or other ligands. In mineral dis- solution, in the detachment of a surface unit into the solution, this metal complex is the precursor of the activated complex in the transition state theory. The stoichiometry and the thermodynamic stability of the surface complexes can be derived from solution based sorption experiments and can be quantified with the aid of adsorption or mass law (surface complexation) equilibria. The mecha- nisms by which metals are released from a dissolv- ing mineral surface are similar to the mechanism of ligand exchange around solute metal complexes and the depolymerization reactions of polymeric metal complexes [5,25,26].

2.3. The crystal suJface

The basic processes which occur during crystal growth or dissolution are sketched in Fig. 2(b). Transport and surface reactions occur normally in series and the slowest step controls the overall rate of mineral dissolution or crystal growth. In this discussion we will concentrate on surface con- trolled reaction kinetics which have been shown to be important for many geochemically important minerals. In crystal growth, growth units (ions, or molecules) become incorporated into the crystal lattice at a growth site, i.e., at a surface step or a kink site. Obviously, ions attached to a step are much more stable than simple adatoms (and ions), because there are more bonds formed with other surface atoms. A loosely adsorbed atom (molecule, ion), undergoing surface diffusion, is commonly termed an adatom. The propagation of steps is basic to orderly crystal growth or dissolution. Tile growth of steps occurs at kink sites. These sites are of special relevance because they self-generate (i.e., an ion or molecule arriving at such a site creates a new similar site). In much of the crystal growth literature dissolution is described as a precipitation process occurring with a negative rate under conditions of undersaturation [27 29]. The surface layers are easily attacked at kink sites. In dissolution the counterpart to adatoms ("ad

146 W. Stumm/Colloids Surfaces A: Physicochem. Eng. Aspects 120 (1997) 143-166

/

_~,/_~

~ M - C

- - M - F

- - M - C

- - M - C / structure solid phase~ oxo, hydroxo complex of M

(a)

H ~ ° ~ J M c ~ e c u l e

3 H

3

• C ~ 3 I

' C ~ 3

Water Molecule

re+

d

ira,

surface complex forma- tion of surface hydroxyl goups with H ÷. Me "+ and ligands

P Inner-sphere complex

outer-sphere m, diffuse io~ complex swarm

9 P electric double layer

Examples of surface complexes:

Positive charge due to surface protonation

possible attachment uf outer sphere surface complex

Inner-spheric (bridging) binuclear surface complex of phosphate

Inner-sphere mononuclear bidentate surface complex of oxalate

Inner-sphere surface complex of F

Negative charge due to deprotonation (reaction with O H )

Possible attachment of outer-sphere surface complex

Inner-sphere monodemate surface complex of Fe(I1)

H H \ /

H/O~ torn

(b)

(c/

Fig. 2. Models on structure and reactivity at the mineral-water interface as exemplified by surface complex formation and atomic surface topography of a crystal. (a) Schematic portrayal of the hydrous oxide surface, showing planes associated with surface hydroxyl groups ("s"), inner-sphere complexes ("a'), outer-sphere complexes ("fl') and the diffuse ion swarm ("d"). In case of an inner-sphere complex with a ligand (e.g., F , HPO42 ) the surface hydroxyl groups are replaced by the ligand (ligand exchange!. (b) Surface topography of a crystal surface; the steps during the growth on the dissolution of a crystal are: (1) transport of species (solvated atoms, ions, molecules) through the parent phase; (2) attachment or detachment of species to the surface; (3) movement of adsorbed species on or into the surface; (4) attachment or detachment of species to edges or kinks. Steps 2-4 are surface processes. (c) Structure of goethite; different coordination conditions prevail on different crystallographic planes (Spadini et al. [18]). At the surface, three different functional groups exist; they are bound to one (A), three (B) or two (C) Fe(III) atoms. The groups B, which are located in the middle of double chains, are non-reactive.

ions") are vacancy defects. In the removal of an ion from a kink site, the kink site is regenerated. As shown by Lasaga [30], at equilibrium the integration rate of ions to kink sites is equal to the detachment (dissolution) rate of kink sites. Precipitation (crystal growth) and dissolution of a solid phase together establish an equilibrium; the principle of microscopic reversibility (detailed bal-

ancing) interconnects these processes. (For equilib- rium to be reached, all elementary processes must have equal forward and reverse rates.) Adsorption of crystal-foreign substrates (impurity ions) especi- ally at reactive sites (such as kink sites) play an essential role in surface dynamics, above all in inhibiting crystal growth or dissolution. An other important defect site that can greatly affect growth

14". Stumm/Colloids Surfaces A: Phvsicochem. Eng. Aspects 120 (1997) 143 166 147

or dissolution, is the screw dislocation at the surface. In screw dislocation growth, the disloca- tion is a source of a step.

The mineral surfaces and particle surfaces are also ordered at the molecular scale, the structure and short range order have become revealed especi- ally by X-ray absorption spectroscopy. Fig. 2(c) exemplifies the structure of goethite on the different crystallographic planes [ 18]. At the interface vari- ous OH groups are bridged to the Fe centers; surface sites of different reactivity can react with H + and OH , ligands and metal ions [16-18] .

All the concepts mentioned complement each other for a full understanding of the chemical interactions at a mineral-water interface and the effects of the structure on reactivity. Information gained from solution chemistry needs to be comple- mented by a structural analysis of the mineral surface to provide information on the structure of the surface complex and to describe the mechanism involved in the adsorption reactions at the micro- scopic scale [18].

2.4. Inner-sphere sur[ace complexes

As Fig. 2{a) illustrates, the functional groups at the surface of a (hydr)oxide or silicate mineral surface react coordinatively with H+, O H , metal ions or ligands to form surface complexes. The extent of surface coordination at the surface sites and its dependence on pH and solution variables can be interpreted in terms of a mass law equation. Surface charge results from the sorption (surface complex formation) reaction itself; the effect of surface charge on sorption (surface complex forma- tion) can be taken into account by applying a correction factor derived from the electric double layer theory to the mass law constants for the surface reactions.

The tendency to form inner-sphere surface com- plexes may be compared with the tendency to form corresponding inner-sphere solute complexes. A few representative free energy relationships (LFER) are exemplified in Fig. 3. Outer-sphere complexes may be formed at the surface if an ion of opposite charge approaches the surface groups within a critical distance; as with solute ion pairs the

"sorbed" ion and the charged functional group are separated by one or more water molecules. Furthermore, charged species may be in the diffuse swarm of the double layer.

in order to understand the reactivity of a mineral surface (the rate of processes such as precipitation and dissolution or catalysis) we need to identify the surface species and its structural identity. The type of linkage to the reactive sites of the crystallo- graphic surface need to be known, above all we need to distinguish between inner-sphere and outer-sphere complexes, between mono-nuclear and bi- or multi-nuclear surface species in order to deduce the precursor of the activated complex for the reaction.

Wieland et al. [34] have postulated that the site energy (Madelung energy} of the most stable lattice constituent, generally a cation at the surface site, characteristic of the free energy needed to break the essential bonds in the lattice, can be correlated with the dissolution rate. Sverjensky and Molling [35] have developed a free energy equation for the prediction of standard free energies of forma- tion of isostructural families of crystalline solids: one term in the relationship accounts for the mineral structure, a second addresses coulombic interactions, and the last term includes ligand field interactions. This relationship can also be used for estimating dissolution rates of isostructural crystal- line solids such as divalent metal oxides and orthosilicates.

3. Dissolution as a ligand exchange reaction

The dissolution (and the formation) of a mineral phase are characterized by a change in the coordi- native environment. Upon dissolution (or crystal formation) the coordinative partner of the crystal constituents change, i.e., in dissolving an oxide or silicate, the oxide or silicate ligands are replaced by H 2 0 or by other ligands. The most important reactants participating in the dissolution of a solid mineral are H20, H + (or H3O+), O H , ligands and in case of reducible or oxidizable minerals, reductants or oxidants. Some of the steps involved in the dissolution of a mineral can be appreciated

148 HL Stumm/Colloids Surfaces A." Physicochem. Eng. Aspects 120 (1997) 143-166

Tendency of a metal ion (trivalent) to form solute complexes with ligands vs tendency to form surface complexes with ligands, e.g.,

MOH 2+ + H2A ~ MHA 2+ + H20 ; KI -=MOH + H2A ~--- -=MHA + H20 ; K]

Data given are for goethite o and i5-A1203 • surfaces. Sigg and Stumm [31] 1t

10

8

6

4

2

0

i i i i f i i i I

]'Phmalic acid / HaPO,,;T? t ~) /

acid C~"ate ch°l /

/ I I i i I I I I I

2 4 6 8 10

log K 1

Tendency of a metal ion to form complexes with OH- (hydrolysis) vs tendency to form surface complexes with oxygen donor atoms, e.g.,

M z+ + H20 ~ - MOH z-I + H + ; *K 1 M z+ + --SOH ~ - -SOM z-I + H + ; *K]

M z+ + 2 H20 ~ M(OH)~ z'/)+ + 2 H + ; *[3 z M z+ + 2 -SOH ~--- (-SO) 2 M (z2) + 2 H + ; *13~

Data are for amorphous silica surfaces. Schindler [32]

~ - I 5 c % ~x

- 10

. , . - 5

_8'

I I 1

p b 2 ~ g2+

J F o . , ,

- 5 - 10 - 15 -20

log "K 1 [*~1

III Tendency of a figand to protonate vs. tendency to bind to surface central metal ion, e.g.,

A 2- + H + ~ - HA ; K21 A 2- +=MOH~ ~ -MA- +H20 ;K~

Data are for hydrous ferric oxide surfaces. Dzombak and Morel [33]

14

12

10

~e~ 8

-~Q 6

4

2

0

' ~ S e O 3 ~

% / , ~ S O 4

I I I I I

2 4 6 8 10

pK2

12

Fig. 3. LFERs between solute complexes and surface complexes on hydrous oxides. (1) Surface complex formation constants are intrinsic, i.e., valid at an uncharged surface.

by c o n s i d e r i n g a m e t a l a q u o c o m p l e x a n d h o w its r eac t iv i ty is in f luenced by l igands a n d p ro tons .

C a s e y a n d c o - w o r k e r s [ 3 6 - 3 8 ] h a v e i l lus t ra ted

tha t the r eac t iv i ty of m e t a l - o x y g e n b o n d s in m i n e r -

als such as ( h y d r ) o x i d e s o r si l icates can be infer red

f r o m the e x c h a n g e k ine t ics o f w a t e r b e t w e e n the

i nne r c o o r d i n a t i o n sphere of the m e t a l ion a n d the

b u l k phase . F a c t o r s t ha t affect the w a t e r e x c h a n g e

ra te of h y d r a t e d ca t ions are the ion ic p o t e n t i a l

(z/r) and in the case of f i r s t - row t r ans i t i on me ta l s

the n u m b e r of d e lec t rons . Because for a g iven

crys ta l s t ruc tu re c o o r d i n a t i o n n u m b e r s and m e t a l

o x y g e n b o n d d i s t ances o f me ta l s in mine ra l s are

s imi la r to these in the h y d r a t e d ion, d i s so lu t ion

ra tes can of ten be co r r e l a t ed wi th wa te r e x c h a n g e

ra tes (Fig. 4).

14A Stumm/Colloids Surfaces A. Phvsicochem. Eng. Aspects 120 (1997) 143 166 149

-10

t - , r ~

~ -12

¢~ -14

-8

Ni

- ~ B e -16 I

3.0 4.0

~Mg

ili Fe

Co

~ M n Ca Zn

I I I 1 5.0 6.0 7.0 8.0 9.0

log k.w(S -x)

Fig. 4. Dissolution rates of end-member orthosilicate at pH 2 and 2 5 C plotted against the first-order rate constant, k w (s ~), for water exchange from the solvent into the hydration sphere of the corresponding dissolved cation. The minerals and ions are identified by the divalent cation (e.g. Mn 2+ refers to the dissolution rate of tephroite (Mn2SiO4) and rate of solvent exchange around Mn(H20) 2~). (From Casey and Westrich [37].)

3.1. Kinetics ~f adsorption, desorption and depolymerization

In the surface controlled formation or dissolu- tion of a solid phase, the adsorption or desorption reaction entails the rate determining step. It will be shown below that the water exchange (rate of exchange of a water molecule between the aquo ions of the metal cation and the bulk H 2 0 ) is a rate determining factor. Consider the adsorption and desorption of a metal ion M' (assumed coordi- nation number = 6) to the surface of a metal oxide

~M OH + M'(H20)6 +

~ M OM'(HiO)~s -" 1)+ + H + ; K ~ (1)

Various authors [ 13,14,39-41 ] have shown that adsorption proceeds first on the formation of an outer-sphere surface complex

= M - O H + M'(H20)~, +

K o 5

"~M O H . . . M'(H20)~ + (2)

which is then followed by a loss of H 2 0 ! +

-=M-OH .. • M ( H 2 0 ) 6

- - , M O . + H , O \ M ' ( H 2 0 ) ~ /n

= - M - O ~ M , ( H 2 0 ) ~ +

fast , ~ M _ O M , ( H 2 0 ) ~ : _ I } + + H"

The adsorption rate depends on the energetics of the interaction as given by the outer-sphere surface complex formation constants, Kos, and, above all, on the adsorption rate constant, kaa.~, which in turn is proportional to the water exchange rate, k-w, of the metal ion (Kos can be calculated with the help of a relation from double layer theory (Gouy Chapman)):

d i e M OM'(H20)~: ~+]

dt t +

= Koskads[M (H20) 6 ] [~M O H ] (3)

Note that the dependence of adsorption on the electrostatic interaction is included in the calcula- tion of Kos. The adsorption rate constant, kads,

can be related to the water exchange rate k w [14,41]; k_w (s 1) measures the rate of exchange of water molecules between the aquo ions of the metal cation complex and the bulk H20. In the adsorption reaction (1) one of the water molecules coordinated to the metal ions has to dissociate in order to form an inner sphere surface complex. Reich and Kohlweit [-42] have suggested that the rate of adsorption for the precipitation of many electrolytes is determined by the removal of a water molecule from the cation during its integ- ration elementary reaction.

3.2. Dissociation

The dissociation (i.e., the desorption) rate is inversely proportional to the stability of the com- plex, K{, (Eq. (1)). Under simplifying assumptions,

d i e M OM'IH20)~-" 1} +]

dt

_ k w K s [ =_M_OM,(H20)~:_ i, + ] [H + ] I4) K I

150 W. Stumm/Colloids Surfaces A: Physicochem. Eng. Aspects 120 (1997) 143-166

where

K~ = K~(int) exp - \ ~ ] (5)

where K~ (int) is an intrinsic surface complex forma- tion constant (valid for an uncharged surface) and 71 is the (electric) surface potential at the adsorption plane of the surface species, F is the Faraday and AZ is the change in the charge of the surface species for the reaction under consideration.

Note that the desorption rate is related to k_ w, the water exchange rate of M', and depends on [H +] and on the extent of the electrostatic interaction.

The rate law for desorption (dissociation) may be looked at as a model for depolymerization (M = M') or dissolution. Indeed the dissociation of

can be characterized by a pseudo first order con- stant k = k t +k2 [ H + ] . One term, ki (s-l) , is for the rate of dissociation into mono-hydroxide species; the other term, k2 (M a s 1), proportional to [H +] corresponds to a pathway where H + ion dependent dissociation (dissolution) reactions have been reported for other oxo-hydroxo multimers, e.g., the decomposition of Alz(OH)24+, (VO)2(OH) 2+ and A11304(OH)~74 - [41,43].

This simple mechanistic view of the H + ion dependent depolymerization may serve as a surrogate model for the proton promoted dissolu- tion of oxide and silicate minerals.

o \ 1 /

Dissolution rate ~ k w

Acid catalyzed

/ on+ - ~ k - -

(6)

H •

/ I o I - - °"*TI - - Dissolution rate ~ kH(C~)"

(7)

It is now well established that the dissolution rate is also related to the amount of adsorbed H +, based on surface speciation modeling and surface titration experiments [6,44-47]. Ab initio quantum mechanical studies on the kinetics and mechanisms of silicate dissolution [48] have shown that the attack of H + (H3 O+) occurs on the bridging oxygen of the Si -O-Si and Si-O A1 structural units on the surfaces of the minerals:

---Si-O-Si~ + H + + H 2 0 ~ ---Si-OH + =-SiOHf (8)

~Si-O A1--- + H + + H 2 O ~ ~SiOH + ~A1OHf (9)

Fig. 4 illustrates the experiments by Casey and Westrich [-37] which exemplify the dependence of dissolution rates of orthosilicates at pH = 2 on the water exchange rate.

The overall rate of the proton-promoted dissolu- tion for aluminium silicates, such as kaolinite has been given by Ganor et al. [49]. Their rate law predicts a linear dependence on the concentration of bound (adsorbed) proton species involved in the reaction.

3.2. Ligand promoted dissociation

The activated complex at the mineral surface resembles the dissolved metal-ligand complex. The introduction of a ligand into the coordination sphere of an aquo metal ion enhances the reactivity of the remaining coordinated H 2 0 molecules (see (10)). Usually the water exchange rate increases with the (r electron-donating, nucleophilic, ability of the coordinated ligand. Simple examples are illustrated in Table 1. As seen in this table, the introduction of an OH ion into a hexa-coordi- nated aquo metal ion enhances the water exchange rate by orders of magnitudes. This reflects the enhancement of the reactivity of the - M - O - bond by deprotonation of the surface functional groups. Thus, as shown in (11) and (12), depolymerization or the dissociation of an M - O - M ' bond, reflecting the dissolution step at the surface of the mineral, is facilitated by O H - bonding or deprotonation or complex formation with a ligand, L.

FV. Stumm/ColloMs Sur/bces A: Physicochem. Eng. A,vwcts 120 (1997) 143 166 151

Table 1

Examples of effect of l igands on water exchange rates of aqua

meta l ions [ 5 2 ]

Aqua meta l ion k

Is ' t

Ca(H20)~, ' 6 x l0 s

Mg(fl2Ol~,+ 3 × 10 ~

FelH_,O)~, ' 4 × 10" FeIHzO);~, ' 1.6 x l02a

F e ( O H ) ( H 2 0 ) { ' 1.2 x l0 s"

Fe(Ot f )2( HzO)4 - 10v (est) b

Fe(OH )a( t 120 /2 ~ 10 + (est) b

AI( H 20)~, + 1 Ni(H20)~ , 3.2 × 10 a N i ( E D T A I : 7 x l0 s

N i ( N H 3)(HeO)~+ 2.5 x l0 s

NiItrien)IH20)_~ ' 5 x Iff' Ni(bipy)( H20)~ ' 4.9 × 104

('r(H20);~, ' 2+4 x 10 ~,a

( ' r ( O H ) ( t l 2 0 ) ~ ~ 1.8 × 10 .+~

where C~ is the surface concentration of negative charge and C{m and C~, are, respectively, the surface concentration of bound O H and of bound ligand L" .

The dissolution rate is then proportional to the concentration of surface bound hydroxide or the extent of surface deprotonation or the concen- tration of surface bound (or adsorbedt ligand [6] .

In both mineral dissolution and in water exchange, the rate of overall reaction is controlled by the rates of breaking of a metal-oxygen bond. Outer-spherically bound ligands have little effect on the water exchange rate [52], and thus have little effect on surface reactivity.

Equilibrium acid-base reactions at surface func- tional groups can account for the pH dependence of the mineral dissolution rates in simple cases.

~' Merbach and A k i n [50] .

h W. Schneider, pers. c o m m u n i c a t i o n [51 ].

I OH° . _ • Bound l igand enhances , + f ' - wa,er exchange ate

- . , , , a v j - - p r emain ing H~O l igands

/ I ( c fTab le 1) -

10)

\ 1 7 / ~--~-" - - OH + ~ M - -

+ / 1 Dissolu t ion rate 7: kouCc'n, or -t k o n C ~

11)

/ ' I 7Y:I .2o - Disso lu t ion rate ~ /q+C',

on+ - '~ l

(121

3.3. Muhidentate ligands

Substitution of more than one water molecule increases still further the rate of exchange of the remaining water molecules (Table 1 I. Thus, coordi- nation with bidentate (Eq. (10it or multidentate ligands give rise to considerable increase in water exchange rate. Thus, bi- or multidentate ligands that form surface complexes with the Lewis acid metal centers of the hydrous oxide or aluminum silicate surface bring electron density into the coor- dination sphere of the surface metal species and labilize the M O- bonds on the surface and enhance the release of metal ions from the surface into the adjacent solution. The labilizing effect of a ligand on the bonds in the surface of the solid oxide phase of the central metal ions with oxygen or OH can also be interpreted in terms of the trans effect, i.e., the influence of the ligand on the strength of the bond that is trans to it, e.g., the effect of a bidcntatc ligand such as a bicarboxylate or the

- - I i o \ /'~1 strength of tile Al-oxygen bonds kl

_ o / \ ~

This influence is attributable to the fact that the ligands trans to each other participate in the orbital of the metal ion; the more a ligand preempts this orbital, the weaker will the bond to the other ligand be [53 ].

152 V~ Stumm/Colloids SurJaces A: Physicochem. Eng. Aspects 120 (1997) 143-166

Ligands which are able to accept electron-den- sity back from the metal ions, on the other hand, such as dipyridyl, terpyridyl and phenantholine tend to have small effects on water exchange rate, as shown for hexacoordinate Ni(H20) 2+ by Margerum et al. [52]. Ludwig et al. [38] report that pyridine ligand, although adsorbed on NiO (bunsenite), do not appreciably affect dissolution rates of NiO.

can form bi- or multidentate mononuclear surface chelates [67]. Many of these ligands introduce negative charge into the surface lattice and increase surface protonation. The rate of ligand promoted dissolution, RL, is proportional to the concen- tration of the ligand at the surface, C~, (mol m -2) [6,34]:

RL = kL C[ (15)

4. R a t e l a w s

As projected by the molecular models, critical metal oxygen bonds are electrophilically attacked by protons which impart surface charge. Protons thus facilitate the release of the metal from the mineral surface. The general proton-promoted dis- solution rate (e.g., m o l m 2 h - l ) may be written as follows [6,34]:

RH = k.(C~)" (13)

where k is the reaction constant, and where C~ is the concentration of surface protons (mol m-2). This concentration (above that of zero proton condition) can be determined by alkalimetric titra- tion. The number n varies between 1 and 3, and often corresponds to the number of protonation steps before detachment [6] , assuming only one mechanism controls dissolution. In the alkaline pH range, there is a basic dissolution due to hydroxide bound on the resulting deprotonated surface groups ( S -OH + O H - H S O - + H20), the corresponding rate law is given by

R o . = kon(Cg . ) m (14)

where C~H is the surface concentration (tool m -2) of deprotonated surface sites (equivalent to bound hydroxo groups) or S O - groups. In case of silicates, C~H is equivalent to C~-sio_ [54]; the number m is usually 1.

Ligands which replace surface hydroxyl groups can attack surface crystalline bonds through a nucleophilic interaction with the metal ions in the surface lattice. Especially effective are ligands whose functional groups contain two or more donor atoms (e.g. di-carboxylic acids, hydroxy carboxylates, diphenols, EDTA, NTA) and which

where kL is the reaction constant. There is a special type of ligand-promoted dis-

solution known as reductive dissolution. This case involves the transfer of electrons from the environ- ment, through a reducing ligand or reductant, to an oxide surface, followed by dissolution. The mechanism of reductive dissolution has been dis- cussed by Zinder et al. [5] and others [55 58], and has been reviewed by Hering and Stumm [59].

In many cases, the rate of reductive dissolution is proportional to the surface concentration of the surface bound reductant (S -OH + HR H S R + H20; charges are omitted for simplicity)

RR = kRC~ (16)

where kR is the constant for the reaction constant and C[ is the surface concentration of the reductant in the surface complex. The proton-promoted, ligand-promoted and reductant-promoted dissolu- tion in terms of plausible simple reaction steps is illustrated in Fig. 5 [ 13].

4.1. Overall dissolution rates: competition by an inhibitor

As indicated in Eqs. (13)-(16), the overall dis- solution rate of a mineral can be described as the sum of the rates of proton-promoted, OH- -p ro - moted, and ligand-promoted dissolution. In case of reducible minerals, a reductant promoted path- way is also possible. The dissolution rate is actually the sum of all the different parallel dissolution rates of the various metal centers. An inhibitor may compete for available surface sites. The dis- solution rate can include the possible effect of an inhibitor; the overall dissolution rate includes the effects or contributions of all surface reactants that

lie Stumm/ColloidY SutJ~wes A: Plo',Yicochem. Eng. Aspects 120 1097) 143-166 153

OH O H organic reduetant

%l~e .~" ~e~(/ ' e.g. ascorbate

/ \ o / xo~ . O y

HO ~ / / 0 -0 fast ~ OH 0

i ~ \ , ~ ' N , d " \ / . , - o / % o f S \ o f e n o / i l ,

OH OH OH O O >< >< l"-- I - e % I-e.,~ I transfer

O H OH / \ 0 / J \ 0 / - C % 0 OH

/ \o. ~.,.., / \o. \%-'~%oZ ~e ''e'/\ ).la.,~e~ OH

Fig. 5. Various pathways for the dissolution of a hydrous oxide exemplified by the proton-promoted (left), ligand-promoted (oxalate as a ligand, in the middle) and reductant promoted (ascorbate as a reductant, right) dissolution of a Fe(lll) (hydr)oxide. In each case, a surface complex (proton-oxalato and ascorbato, respectively) is formed, which assists in the breaking of the -Fe-O(OH) bonds so that a Fe(III) aquo or Fe(IIl) oxalato complex or in case of the reductive dissolution on Fe( It ) complex detaches into solution. In the redox reaction of Fe(III) with ascorbate, the ascorbate is oxidized to the ascorbate radical A--. In each case the original surface structure is reconstituted, so that the dissolution can continue to proceed in maintaining a steady state condition. The surface complex is the precursor to the activated complex; the dissolution rate is proportional to the concentration of the surface complex. The principle of proton-promoted and ligand- promoted dissolution is also valid for the dissolution (weather- ing) of silicates and aluminum silicates.

become sorbed ( b o u n d ) to the meta l centers of the minera l surface [58] .

Riot = ku(C~)" + ~ kL,( u + ~ k~,C~, + ~ kRiCRi i i i

+ koH(C~oH) " + ~ kMiC~i + kaq (17) i

The terms in this equa t ion represent the contr i - bu t ion made to the d isso lu t ion rate by each surface site covered with a l igand, E~kuC[, an inhibi tor , E~k,C~, a metal ion, E~ku~C~a~, or a reductant ,

Ei kRiC~Rs. The terms ku(Ch) ' , and koH(C~H)', repre- sent acid and basic dissolut ion, respectively, The last term in Eq. (17) , kaq, is due to the effect of hydra t ion , and it reflects the po r t i on of the dissolu- t ion rate and usual ly occurs at p H values near the zero poin t of charge. Each of these terms in Eq. (17)

can be in terpre ted in terms of the ac t iva ted com-

plex theory. In this case, for each dissolut ion pa thway each type of metal center complex is considered to be a precursor to the " 'activated" complex. The " 'activated" complex itself can then be released, t ak ing with it the surface metal ion (or its complex), into solut ion. The rate of reductive d issolut ion of an F e I l l l ) o x i d e is given by RRed :

LR C~ Fe(III)R" Since some surface sites (e.g., kinks, steps, defect

sites) are very reactive while others are not, all the individual k terms in Eq.{17) include a mole fract ion x~ of reactive sites [34] . For reasons of simplicity, the model out l ined here assumes that the surface area is constant , that the d issolut ion rate is cont ro l led by surface reactions, and that back react ions such as nucleat ion and crystal growth are negligible. This implies that the active sites are con t inuous ly regenerated dur ing dissolu- tion. Thus by main ta in ing a cons tant rat io (Xi) Of active sites to total sites, a s teady state condi t ion is obta ined . Al though the definit ion of v~ is not l imited to chemical ly and s t ructura l ly equivalent surface hydroxyl groups, the theory of mean field stat ist ics (e.g., [10]) al lows these groups to be t reated as such.

This model , m agreement with exper imenta l results, also predicts that d issolut ion is l inear with time. Since all the dissolut ion react ions are parallel , usual ly only one or two mechanisms d o m i n a t e the react ion under given condi t ions . As a result, of all the terms in Eel.(17), usually only one or two will be of impor tance .

4.2. SutJhce complex Ibrmation equilihria

Because the fo rmat ion of a surface complex is fast c o m p a r e d to its de tachment into solution, it is reasonable to assume that equi l ibr ium condi t ions , with respect to surface complex format ion, are ma in ta ined at all times. The var ious surface equi- l ibr ium concent ra t ions , C~, in Eq. (171, (extent of surface p ro tona t ion , surface concen t ra t ion of l igands or inhibi t ing species) can be exper imenta l ly de te rmined from a lka l imet r ic t i t ra t ion curves and from analyt ica l measurements on the extent of a d s o r p t i o n These concen t ra t ions may also

154 IJd Stumm/Colloids Surfaces A." Physicochem. Eng. Aspects 120 (1997) 143 166

be obtained mathematically from equilibrium constants.

Once a set of surface complex formation equilib- rium constants is established, mass law equations can be used to quantitatively describe the various sorption reactions with H +, OH , metal ions, and ligands. The surface charge is a result of the surface reactions themselves, and its effect on sorption may then be taken into account using the electric double layer theory (e.g., Gouy Chapman model for the diffuse double layer model, or the constant capacitance model). Such calculations and their procedures have been demonstrated by various authors [ 10,11,13,33], and suitable computer pro- grams are available (Westall, 1980). For such calcu- lations, for a given surface area and a given concentration of hydrous oxide, the following information must be known: total surface concen- tration of metal centers (site density), total concen- trations of adsorbable species (metal ions, ligands, inhibitors, etc.), pH (assumed to be constant), ionic strength, and the correspondence surface complex formation equilibria.

5. Specific dissolution rates

In Fig. 1 we illustrate the pH-dependence of the dissolution rate of a few selected minerals. The pH dependence reflects the sum of Eqs. (13) and (14).

In Table 2 representative examples of specific dissolution rates (i) under low pH conditions, (ii) under alkaline pH, (iii) of ligand promoted dissolution, and (iv) of reductive dissolution are given.

To facilitate comparison, the rates are expressed as first order rate constants (time 1). In case of ligand-promoted or reductant-promoted dissolu- tion, the measured rate (mol m -2 h -1) is divided by the concentration of the ligand or reductant surface complex (molto-z). In case of proton- promoted dissolution, where rates may not be linearily related to the proton surface complex (C~), the dissolution rate, Rn, (mol m - 2 h 1) is divided by the total concentration of surface sites, Sr (molto-2), so that k (h-1)=RH/S~. [41]. For all minerals given in the Table, S t = 2 x 10 -5 m o l m -2 (12 sites per nm -2) is assumed.

Different methods used by different investigators (measurements of steady state rates vs. rates from batch experiments; different analytical procedures to determine C~, C[ and C~; different ionic strength) make strict comparisons difficult. As has been shown already by Wehrli et al. [41] the dissolution of A1 minerals at low pH (compare nos. 7-13 in Table 2) show remarkably different rates depending on the crystal habit. Corundum (sap- phire) is the most resistant towards acid dissolu- tion. The faster dissolution rate of (~-A1203 compared to corundum is to be expected 1-41] since this spinel-type solid contains randomly dis- tributed tetrahedral sites. Bayerite dissolves fastest because of its layered two-dimensional structure. The presence of silicate in muscovite and kaolinite reduces the acid AI(III) dissolution (in comparison to 6-A1203, by nearly an order of magnitude; the critical point of attack by the proton is the oxygen atom that bridges a Si and A1 atom [59]; the detachment of AI from the lattice structure is probably the rate determining step. The detach- ment of an Fe(IIl) center from a geothite surface under acid conditions is considerably more difficult than the detachment of A1 from its oxides. Under comparable conditions BeO dissolves faster than 6-A1203 .

In ligand-promoted dissolution, the bi- or multi- dentate ligand that coordinates in mononuclear inner-sphere surface complexes is most efficient in labilizing the =A1-O bond, detaching the surface metal centers into solution. The ring size of the surface chelate and the number of donor atoms which ligate to one surface metal center are impor- tant. As shown by Furrer and Stumm [6], the five-membered surface chelate ring of oxalate is more efficient than the six-membered chelate ring of malonate and salicylate and the seven-membered chelate ring of succinate and phthalate in enhanc- ing the dissolution rate of Al-minerals (cf. nos. 1 10, 12, 13, Table 2). Monodentate organic sur- face complexes have little effect on the dissolution of 6-A1203. Complex formers generally form, if at all, relatively weak surface complexes on silica surfaces; but the nucleophilic citrate and oxalate enhance the dissolution rate of quartz [69], although experimental evidence for the adsorption of these chelates on the surface of quartz is lacking.

W Stumm/Colloids Surfaces A: Physicochem. Eng. Aspects 120 19971 143 166 155

Table 2 Examples on specific dissolution rates

Mineral Surface group Type condition k ( 10 -~ h ( 2 0 C 25 C)

Reference

1 ,~;-A1203 -AI oxalate bidentate mononuclear 2 ,$-AI203 AI malonate bidentate mononuclear 3 &A1203 AI salicylate bidentate mononuclear 4 &AI203 AI succinate bidentate mononuclear 5 ~)-AI203 AI phthalate bidentate mononuclear 6 ~-Al203 AI benzoate monodentate mononuclear 7 ,5-A120: A1OH~ ~ pH = 3 8 Corundum AIOH~ ~ pH = 3 9 Muscovite AIOH~ ~ pH = 3

1(1 Kaolinite : A1OH2 " pH = 3 11 Bayerite AIOH2 + " pH = 3 12 Kaolinite :AI oxalate 13 Kaolinite AI salicylate 14 BeO BeOH + a pH = 3 15 ~-F'eOOH FeOH + ~ p H = 3 16 :~-FeOOH Fe oxalate pH = 3 17 ~-FeOOH ICe oxalate pH - 5 18 ~-FeOOH Fe oxalate p H = 5 19 y:l:eOOH Fe EDTA 3 ,?a pH = 5 20 ~-FeOOH Fe EDTA 3 mononuclear pH = 8 21 x-FeOOH Fe EDTA 3 binuclear ¢ pH = 5 22 :~-t'e20 ~ Fe oxalate pH = 3 23 >t:e2().~ Fe oxalate pH = 3 24 x-t:e203 Fe oxalate pH = 5 25 ~-t"e203 :l::e EDTA 3 pH = 5 26 :~-l:e203 : Fe ascorbate reductive pH = 3 27 ~-F'e,O~ Fe ascorbate reductive pH - 4 28 ):-]:e,() 3 FeS sulfide reductive pH - 5 9 29 ~-I:ee().~ FellS sulfide reductive ptt = 3 9 30 i'-FeOOH Fe EDTA 3 pit - 3 10 31 i,-t:eOOH Fe malonate pH = 3 32 Quartz : SiOH ~ pH - 3 33 Quartz SiO pH > 5 34 Kaolinite A I O H + " pH = 3 35 Kaolinite SiO pH > 7 36 Albile SiO pH > 7 37 Calcite CO3 H " pH = 8

I 1 Furrer and Slumm 116] 7 Furrer and Stumrn [6]

13 Furrer and Stumm [6] 2 Furrer and Smmm [6] 3 Furrer and Stumm [6]

< I Furrer and Stumm [6] 0.6 Furrer and Stumm [6] 0.44 Caroll-Webb and Walther [45] 0.07 Wieland and Slumln [59] 0.15 Wieland and Stulmn [59]

27 Puller et al. [63] 1.7 Wieland and Stumm [59] 2.4 Wieland and Stumm [59] 2.2 Furrer and Stmnm [6] 0.05 Zinder el al. [75]

12 Zinder el al. [5] 3 Zinder et aL [5] 0.7 Borgaard, 1991 0.7 Borgaard, 1991

68 Nowack and Sigg [65] 1.6 Nowack and Sigg [65]

35 Zinder et al. [ 5] 13 Zinder ct al. [ 5 ]

1.4 Borgaard [64] 1.0 Borgaard [641]

1 I 0 Sutcr ct al. [ 57] 45 Surer et al, [57] 30 Dos Sanots and Stumm [58]

400 Don Santos and Stumm [58] 200 Bondietti et al. [66]

20 Bondicni et al. [66] 0.11 Brady and Walther [67]

11.3 Brady and Walther [67] 0.21 Wieland and Stumm [59] 0.6 Caroll and Wahher [68]

11.3 Brady and Walther [68] 105 Chou and Wollast [ I ]

AssumingH4 a n d S r = 2 X 10 ~molsur faces i t esm 2, k_Ru,S. t . b Probably binuclear bridging. c Predominantly binuclear.

EDTA can participate with up to six donor atoms in forming surface chelates with hydrous oxides of Fe( l lI ) and AI(III). Its effect on surface reactivity appears to be critically dependent on whether a mononuclear or polynuclear surface complex is formed. In case of goethite, N o w a c k and Sigg [65] have experimentally established that

in case of goethite below pH = 7 binuclear or multinuclear EDTA surface complexes are formed, while above pH 7 mononuclear surface chelates are formed. The enhancement of the dissolution rate in this case is very remarkable. At pH 5 (bi- or multinuclear surface complex) the dissolution rate per adsorbed EDTA is forty times slower but

156 [4~ Stumm/Colloids Surfaces A: Physicochem. Eng. Awects 120 (1997) 143-166

Table 3 Effect of various surface interactions on enhancing or inhibiting the surface controlled dissolution (modified from Biber et al. [60])

pH adjustment

Ligand adsorption

Multivalent metal ion adsorption

Ionic strength

Oxoanions (such as phosphate, silicate, arsenate, borate, chromate, sulfate)

Surface polymerization and surface precipitation

Reductants

Isomorphic substitution (formation of solid solutions) at surface

Affects surface protonation and deprotonation, The dissolution rate is usually minimized at pHzp c. Dissolution rate is enhanced by inner-sphere ,nononuclear surface complex formation. (Nucleophilic interaction with central metal ion.) In addition, adsorption of ligands, that form negatively charged surface complexes, enhances surface protonation. This, in turn, may increase dissolution rate. Usually decreases the dissolution rate because, ( 1 ) reactive surface sites are blocked from interaction with ligands and protons, and (2) there is a decrease in surface protonation. Binuclear inner-sphere surface complexes very efficiently inhibit acid dissolution. On the other hand, basic dissolution is accelerated because of surface deprotonation. Changes surface speciation because of double layer effects. Increasing I increases positive and negative surface charge at pH below and above pHzec, respectively. Thus, it increases dissolution rate. Tend to form bi- or multinuclear inner-sphere surface complexes which are dissolution-inert and occupy otherwise reactive sites. Bi- or tri-nuclear surface complexes and bi- or trivalent ligands, that form uncharged surface complexes, do increase surface protonation and are most efficient passivators. Adsorbed cations or anions may form polymers which inhibit dissolution by crosslinking the surface lattice; at high surface coverage, a surface precipitate (film) may form with the constituent ions of the mineral. Such a film acts as a protective layer. Adsorption of reductants to reducible oxides or semiconducting oxides facilitates subsequent electron transfer and favors reductive dissolution. Oxidants may act as inhibitors because they prevent reductive dissolution. Changes Lewis acidity and site energy, as well as kinetics of dissolution.

most l ikely includes also the effect of a small fract ion of m o n o n u c l e a r complexes. Wi th lepi- docroc i te Bondie t t i et al. [ 66 ] found in the p H - r a n g e 3 -10 m o n o n u c l e a r EDTA and ma lona t e surface complexes. The enhancement of EDTA on Fe ( I I I ) d issolut ion is within a factor of 2.5, s imilar to that found with goethite.

The b inding of an add i t iona l O H - ion in the surface complex or the d e p r o t o n a t i o n of a surface O H - g roup ( S - O H + O H ~ S O - + H20}, i.e., the in t roduc t ion of a negat ive surface charge at a metal center, facili tates d issolut ion in the a lkal ine pH-range . In the a lkal ine pH range the d isso lu t ion of SiO2 and of silicates has been shown to depend on the surface concen t ra t ion of [---SiO ] [67] .

6. Inhibition vs. dissolution

The rate of d issolut ion of a given hydrous oxide is a lways enhanced or inhibi ted with respect to a

reference rate [60] . The inhibi t ion of this dissolu- t ion rate is de te rmined by the fol lowing majo r factors: ( 1 ) the surface concen t ra t ion of the relative inhibi tor , C~ of Eq. (17), and (2) the kli of Eq. (17). The surface densi ty of the relative inh ib i tor reflects how well it adsorbs in compe t i t i on with o ther adsorbates . It is i m p o r t a n t to note that the adsorp - t ion of a species such as an inh ib i tor also affects the surface concen t ra t ions of o ther surface species. (Cat ion and anion adsorp t ion , for example, are accompan ied by decreases and increases, respec- tively, in surface p r o t o n a t i o n (C~), due to con- s t ra ints of the electric double layer.) Whi le the surface concen t ra t ion of inhib i tor should be high for op t imal inhibi t ion, the value of kii in Eq, (171 must be small (with respect to ku, k t , etc.) or even zero. If the surface concen t ra t ion of the inh ib i tor is high, then there are less avai lable sites for d i sso lu t ion-enhanc ing l igands or p ro tons to adsorb . If the value of k~ is then small, the overal l rate given by Eq. (17) will be low, co r re spond ing

if'. Stumm/Colk~ids Surlaces A: Phvsicochem. En~. Aspects 12# (1997) 143 166 157

to a reduction in the dissolution rate. The various k~ values in Eq. (17) depend on the probability of finding a specific site type in a geometric and coordinative arrangement conducive to the forma- tion of a precursor complex. This precursor com- plex may subsequently detach into solution. The k~ wflues also depend on the values of xi, the mole fraction of active sites of type i. An inhibitor can thus reduce the k~ values by either producing surface configurations which inhibit or prevent dissolution, or by reducing x~, e.g., by preferentially adsorbing to different surface sites such as kinks or steps.

Ligands capable of forming bi- or multinuclear surface complexes, i.e., ligands that can bridge two or more metal centers from the surface lattice, essentially extend the cross-linking of the "'poly- mer" or solid. They are believed to be inhibitors because the simultaneous removal of two or more metal centers is energetically unfavorable. As will be shown below, oxoanions such as phosphate, arsenate, borate and silicate are most likely to form such bridging surface complexes, above all at kink or edge sites. Additionally, the complexes formed by oxoanions in solution are characterized by low stability constants, indicating that the oxoanions form rather unsuitable leaving groups (i.e. activated complexes) for dissolution. Specifically adsorbed cations can block surface groups and enhance surface deprotonation; they are surface poisons for crystal growth and dissolution. Furthermore, cat- ions often form binuclear surface complexes [16 18]. Some cations such as Cr( l l l ) and possi- bly also AI(III) form polynuclear clusters on the surface.

7. A few experimental case studies on the inhibition of oxide dissolution

Fig. 6 illustrates a few representative results on effects of inhibitors on ligand promoted non-reduc- tive and reductive dissolution (Figs. 6(a), (b)). EDTA is a powerful ligand in promoting the dissolution of Fe(lll)(hydr)oxides. In case of lepi- docrocite, EDTA is assumed 1-66] to form mono-

nuclear surface complexes, i.e., several of its donor atoms ligate to one Fe(III) surface center. The pH dependence of EDTA-promoted dissolution is explainable in terms of the pH-dependence of EDTA adsorption. (For EDTA-promoted dissolu- tion of goethite see Nowack and Sigg [65].) As shown in Fig. 6(b), arsenate, phosphate, selenite and to a lesser extent SO]- are powerful inhibitors at neutral and slightly acid pH wflues. These inhibitors most likely compete successfully with EDTA for suitable reactive surface sites. At low pH (pH<~3) lhe oxoanions, phosphate, selenite and arsenate accelerate the dissolution of ;,-FeOOH in presence of EDTA. A corollary to these findings is the observation that the corrosion rate of passive iron in nitric acid is markedly accelerated by traces of H3PO4, while phosphate in the more neutral pH is a most valuable corrosion inhibitor [70]. A hypothesis advanced by Bondietti et al. [66] assumes that these oxoanions form at low pH primarily mononuclear surface complexes, while around pH = 7, primarily bmuclcar com- plexes are formed. This hypothesis finds support by studies on surface complex formation equilibria which finds this pH dependence on the nuclearity of the surface complexes for selenite 171] and phosphate [31 ].

The reductive dissolution of Fe(111) minerals by a reductant such as H2S is much faster than ligand- or proton-promoted dissolution. The dissolution reaction, as shown by Dos Santos-Afonso and Stumm [58], is initiated by the formation of I-:oS and ~FeHS surface complexes: the subsequent electron transler within the complex leads to the formation of Fe(ll) centers in the surface lattice. The Fe(ll) O bond is characterized by a smaller Madelung energy than the Fe(II1) O bond (larger radius of Fe(lI) and larger H 2 0 exchange rate than Fe( lll)l: therefore the Fe(ii) is more readily detached from the surface into the solution.

As shown in Fig. 6(d), phosphate and borate inhibit the dissolution of goethite by H2S. Similarly, the dissolution of lepidocrocite (;,-FeOOH) by EDTA, (Y4) is inhibited by phos- phate and arsenate ( Fig. 6(a)). Both in the reduclive dissolution {by H2S) and the ligand promoted dissolution (by EDTA) the inhibition effects can be explained by ligand exchange reactions of the

158 W. Stumm/Colloids Surfaces A." Physieochem. Eng. Aspects 120 (1997) 143-166

1.5×10- 4

1 × 1 0 -4

5×10- 5

a) 0.5 g/g lepidocrocite ~ / 10 -3 M EDTA H " 0.01 M NaCIO, ~ p H S ~

5 10 15

time [hours]

tool m -2 h l

2 ~ 10 -7

10-7

b) , ~ 0.5 g/g lepidocrocite

0.01 M NaCIO 4 ] 10 "3 M EDTA

C) l0 "3 M EDTA onl Q ® s~ ,O-~M

~ ] ~phl2;;h:t 'e 0;;-tMM

~ [ ~ Arsenate 10-3M _

® ®® ~~ 1 I I I I

3 5 7 9 II

pH

c)

70 60 ~ I H 2 Sal°ne 0.01MNaC104

50 I- q o.o17 g/~Hematite

'Z'%- 40 f 4 p H : 6

52 30 I-4 101.tM ,

I0

0

0 2 4 6 time (h)

10 7 --

10-9 L

10-n -

k-

10-13

lO-IO

2 x 10 -4

10-4

r~ ~3

0

d) 1 I I t I . "

¢• 0.001 arm H2S H2S only , " " 0.01 M NaCIO 4 oH=5 •¢ 0.032 g/g goethite , - " s

• . I " "

• , 0.1 M borate

..~ ..4 ~ ~ 0.001 M phosphate

100 200 300

time [mini

• / /

0.001 ann H2S

0.01 M NaCIO 4 p H = 5 0.032 g/g goethite

I I I

10-8

I

10-6

e) [ I I I I 1 I I I

---FePO4H" x . ~ ~ . . . . . . . o- . . . . . . .13 . . . . . . . I:]- . . . . . - -

-=FeS H

I I I I

0.0001 0.01

total phosphate [M]

w. Stumm/ColloMs Surfaces A: Physicochem. Eng. A.wects 120 / 1997) 143 166 159

type [the definitive assignment of the pro tona t ion of surface species has not yet been established).

2 -FeS + H 2 P O 4 + 3H + = (=Fe)eHPO4 + 2H2S (18)

2 - F e H 2 Y + H 2 P O 4 + H +

= ( Fe)2HPO4 + 2H3Y (191

The fact that oxoanions can effectively inhibit reductive and non-reductive dissolution with regard to a reference system, supports Eqs. (18), (19) and the concept of competi t ion between dissolut ion-promoting and dissolution-inhibiting ligands. As Fig. 6(e) shows with the help of M I C R O Q L calculations [62] ; based on complex format ion constants fitted to actual data by Stigg and Stumm [31] on the one hand and by Dos Santos-Afonso and Stumm [58] , on the other hand, the dissolution system is very sensitive to phosphate [60] .

Fig. 7(a) gives an example of an inhibition of the reductive dissolution of Fe30 4 by a polyelectrolyte (polyacrylic acid). The inhibitory effect has been assumed [72 ] to be the result of the competi t ion between the polyelectrolyte and the reactant mole- cules for the active sites in the oxide particle. The inhibitory effect of cations is illustrated in Fig. 71c) which shows, that Cr(I I I) is a very efficient inhibi- tor of hematite (~-Fe~O3) dissolution. The adsorp- tion of Cr ( l l I ) to goethite surfaces has been investigated, using X-ray adsorpt ion spectroscopy, by Charlet and Manceau [ 16]. Scanning tunneling microscopy was used by Eggleston to study the adsorpt ion o f C r ( I l I ) to hematite (001) [73] . These studies confirm inner-sphere adsorption, and sug- gest that the most stable sites for Cr ( I I l ) on the surfaces are associated with octahedral vacancies in the underlying oxide and, depending on surface density, show a preponderance of adsorbate clus- tering. Cat ions exhibiting a strong affinity for

mineral surfaces such a s g o 2+ or AI 3+ may inhibit the dissolution of A1203, of albite and of mon tmo- rillonite (Fig. 8(c)).

Very thin oxide films (passive films) have been shown to have a p ronounced influence on the corrosion rates of metals and alloys; the factors that control the dissolution behavior of the most important oxide phases, that constitute the passive films on carbon-based and stainless steels ( i ron(I l l ) , ch romium(I l l ) and mixed oxides), are of relevance in corrosion protection and in the chemical decontaminat ion of sled surfaces [74 77].

8. Mechanisms involved in adsorption reactions at the microscopic structural scale

As we have seen, dissolution or inhibition is initiated by sorption reactions at suitable surface sites. Detailed information has been gained, above all from EXAFS spectroscopy, on the nature of surface complexes on ferric hydrous oxides, especi- ally goethite and ferrihydrite [15 18,78 80 I . Obviously, the simplified model, the surface as an array of equivalent surface hydroxyls, needs refinement to distinguish surface sites of different reactivity. Not only do we need to differentiate between higher affinity and lower affinity sites of different linkage mechanisms (at differently unsatu- rated surface sites), involving one or two bonds (mono- or binuclear surface complexes) but hax.e to consider that different reactivity groups for surface complexing may occur in different crystal- lographic planes (e.g., Fig. 2(c)1.

In crystals the area of crystallographic planes is inversely related to their reactivity. As pointed out by Spadini et al. [ 18 ] the growth rate is highest in the direction perpendicular to small planes which, correspondingly, hold the sites of high

Fig. 6. Examples of ligand promoted and reductive dissolutions of Fe(llll(hydr)oxides and their inhibition. {a) Non-reductive dissolution of lepidocrocite by 10 3 M EDTA at various pH values. Data: Bondieni et al. [66]. (bl Rates of lepidocrocitc dissolution by 10 .1 M EDTA vs. pH and in presence of 10 2 M SO4 , 10 -~ M selenite. 10 3 M phosphate. 10 .t M arsenate, respectively. Data: Bondietti et al. [66]. (c) Reductive dissolution of hematite by H2S {0,01 atm) at p t t - 6 and in presence of 10 ~ M al~d 10 a M phosphate, respectively. From Biber et al. [60]. {d) Reductive dissolution of goethite by H2S ( 10 "~ atml at pH = 5 and m presence of 0.1 M borate and 10 3M phosphate, respectively. From Biber et al. [60]. le) Sur|'ace equilibrium compctilion bel~een reductant (lt2S) and inhibitor (phosphate). Modified from Biber et al. [60 ].

160 W. Stumm/Colloids Surfilces A." Physicochem. Eng. Aspects 120 (1997) 143-166

4

3

2

1

0 0

a)

] ~ 0.267 g~g Fe304

/ / o \ ~) 0. ' 47 ×_; 03 Ml)~aaic°gllYC°cfid acid / 0y t I 1 I r I

2 3 4 5 6 7 8

pH

r,

b) 8

6 '

4

2

0

without C r ( l l I ) ~

J • 0.5 g/g Goethit (k01 M KNO 3

J prl=3

0.001 M CrOlI)

20 40 60 80

time [hoursl

c)

-7.5 I ~ Albite

~ -8.0 -

o

~ -8.5 - N

-9.0 I I I

-7 ~5 -5 ~l -3

log [AI(III)I [mol/g l

Fig. 7. Examples of dissolution and inhibition. (a) O Specific rate constants for the reductive dissolution of Fe~O4 by thioglycolic acid and O inhibition of thioglycolic acid mediated dissolution by polymethacrylic acid. (From Baumgartner et al. [72].) (b) Dissolution of hematite at pH = 3 and in presence of 10 3 M Cr(III). (From Bondietti et al. [66].) (c) Dissolution of the feldspar albite at pH - 3; the logarithm of dissolution rate is plotted as a function of log[-Al(IIl)] in solution• Data from Chou and Wollast [ 1 ]. For modeling the inhibition effect of AI(III) see also Wehrli et al. [41 ].

affinity. In ~-FeOOH (Fig. 2(c)) the affinity of surface sites for cations decreases in the order A(ool) > A¢olo) ~> A~11o} ~ A(lo0). Cations, such as Cd ions, are expected to have the same relative affinity for crystallographic planes of ~-FeOOH as do Fe 3 + ions during crystal growth (i.e., most of the possible Cd sorption sites correspond to crystal growth sites). This insight has consequences for dissolution and its inhibition and on crystal growth. One can infer that dissolution occurs at the planes and sites of highest reactivity. Thus, dissolution may occur preferentially perpendicular to small planes.

Organic ligands with two or more donor atoms such as oxalate, malonate, citrate, phthalate, salicy- late, and EDTA are sufficiently flexible to form mononuclear surface complexes and thus transfer electron density most efficiently to the surface metal centers assisting in the disruption of bonds and their detachment. Inhibiting species most likely interact also preferentially with the functional sur- face sites of high reactivity and thus block the critical reaction sites. Blocking of the same sites, under conditions of oversaturation, interfere with crystal growth.

The local structure of hydrous ferric oxide (HFO) has strong similarities with 7-FeOOH and differs from that of ~-FeOOH (lepidocrocite) with respect to the relative distribution of O and O H groups around Fe atoms [-18].

8.1. Oxoanions

Anions such as phosphate, arsenate, selenite, borate, silicate, carbonate, and sulfate are believed to sorb on goethite and other hydrous ferric oxides primarily by exchanging two A-type surface hydroxyl ions (binuclear surface complexes briding two corners of the polyhedra) (Fig. 2(c)). These bridging surface complexes extend the cross-linking at the surface lattice; they are also believed to be inhibitors of dissolution because the simultaneous removal of two or more metal centers is energeti- cally unfavorable.

8.2. Surface poisoning

As Fig. 8 illustrates, arsenate (and phosphate and other oxoanions) attach in grooves of the

l,t~ Slumm/Colloids SurJbces A." Physicochem. En,e. Aspects 120 r 1997) 143 166 161

Fig. 8. Idealized geometries of arsenate surface complexes. Short dioctahedral Fe oxyhydroxyl chain saturated with tetrahedral (bridging binuclear) arsenate complexes; perspective drawing of a small unit of goethite structure with adsorbed arsenate on the groove surface: complexed arsenate, similar to phosphale and most likely' other oxoanions, occupies the positions where ferric dioelahedral chains are linked by octahedral corners to adjacent chains. Thus, adsorbed arsenate, and mosl likely other oxoanions, could serve to disrupt further crystallization of lhe goethite surface. Modified from Waychunas et al. [78].

goethite surface parallel to the c-axis [71,82]. The grooves correspond to the edges of the double ferric octahedral chains that make up the goethite, and are present in the akageneite, lepidocrocite and ferrihydrite structures. Thus, adsorbed arse- nate (and phosphate) could not only block critical surface sites for dissolution but could disrupt fur- ther crystallization of goethite. Similar inferences can be drawn for the growth inhibition of other Fe(l l l ) lhydrloxides by arsenate and phosphate [78]. On the basis of experimental data, Waychtmas et al. [78] assume that the precipita- tion of ferrihydrite occurs in the following manner: isolated ferric hexacoordinated species link edges to form small chains of Fe-octahedra. In some planes two chain edges can be shared, leading to the development of small plate-like species of dioctahedral and trioctahedral chains, This process continues until chain species have reached suffi- cient concentration that cross chain linking pro- cesses by octahedral corners is the favored process. The presence of arsenate (or phosphate) during the stage of crystal growth stops the Fe -O Fe chain- linking process and retards the polymerization, even at relatively moderate levels of arsenate load- ing. In classical crystal growth this would be termed surface poisoning. It should be noted that

the mechanism of arsenate adsorption described does not constitute surface precipitation or solid solution formation [78].

The bidentate binuclear bonding of arsenate to neighboring iron octahedra has been further sup- ported by iron isotopic exchange and M6ssbauer spectroscopy oil the ferrihydrite surface [80]. The kinetics of iron isotopic exchange with Fe-NTA solutions showed that there are at least two types of iron sites. One population of sites (the labile ones) approached iron isotopic equilibrium within 24 h while the second population of sites (the non- labile ones) exhibited a much slower rate of isotopic exchange. Adsorbed arsenate reduced the degree of exchange by labile sites, indicating that the oxoanion blocked or greatly inhibited the rate of exchange of these sites. The labile population of sites has been assumed [80] to be composed probably of end sites of the dioctahedral chain structure of ferrihydrite.

These studies support also earlier investigations and provide additional interpretation on the effect of suitable ligands on the transformations of ferry- hydrite into goethite and hematite [83,84]. Silicate and phosphate, for example, stabilize ferrihydrite and hinder the formation of goethite (which occurs upon dissolution and subsequent reprecipitation) and thus, indirectly favor the formation of hematite (which is formed by internal aggregation and rearrangement I.

&3. Microtopographic observation

Atomic force microscopic (AFM) observation on the dissolution of hematite [85] confirmed that dissolution manifest itself preferentially at high- energy sites (dislocation steps, edges, kinks). Dissolution by citric and oxalic acid (pH = 3) occurred primarily by the retreat of step, edges and the formation of edge pits: the experiments demonstrated that the bulk of the basal-plane surface of hematite is unreactive with respect to dissolution. These observations are in agreement with the known effects of citrate on hematite crystal growth. Schwertmann et al. [86,87] demonstrated that citrate alters the shape of hematite particles grown from solution and attributed these differ- ences to pre[erential citrate adsorption on the

162 V~ Stumm/Colloids Surfaces A: Physicochem. Eng. Aspects 120 (1997) 143-166

prismatic planes. Studies of phosphate sorption also have shown that the basal plane surface of hematite is also unreactive compared with the prismatic surfaces [87,88]. They attributed the increased sorption of phosphate on the non-basal plane surfaces to the A-corner hydroxyls on the non-basal (110), (100) and (223) surfaces and the absence of monocoordinated A-hydroxyls on the basal (001) surface.

Theories of dislocation-related etch pit nucle- ation have been reviewed by Blum and Lasaga [44]. Fig. 9 illustrates some of our observations [89] on the effect of acid deposition on granitic gneiss. The surface of some white, macroscopic plagioclase grains were covered with etch pits. The presence of etch pits suggests that defects in the crystal structure are sites of strong preferential dissolution during geochemical processes [44]. Such etch pits are also evidence for surface-con- trolled (rather than diffusion or transport-con- trolled) dissolution reactions. As shown, the etch pits were arranged along subgrain boundaries.

8.4. Metal cations as inhibitors

Sorbed cations, such as Cu(II), g o 2+, Co 2+, Ni 2+ and Cd 2+ have been shown on many oxides to be sorbed as inner-sphere binuclear surface com- plexes. These ions are thus potential inhibitors for crystal growth and dissolution. Spadini et al. [ 18] have shown that at the most reactive (001) plane of goethite Cd (octahedra) adsorbs primarily by sharing two edges and two corners of Fe(III) octa- hedra; this is binuclear binding. At the (100) and (110) plane a Cd octahedron can above all establish a two corner linkage by bridging two adjacent A positions; these are sites that block the three-dimen- sional condensation of octahedral chains along their length [16]. The adsorption of Cr(III) to goethite surface has been investigated using X-ray adsorption spectroscopy by Charlet and Manceau [ 16]. Scanning tunnelling microscopy was used by Eggleston to study the adsorption of Cr(III) to hematite (001) [73]. These studies confirm inner- sphere adsorption, and suggest that the most stable sites for Cr(III) on the surfaces are associated with octahedral vacancies in the underlying oxide and, depending on surface density, show a preponder- ance of adsorbate clustering.

8.5. Surface polymerization and surface precipitation

Fig. 9. Scanning electron micrograph of weathered gneiss, showing packs of mica splitting into sheets. Attack was apparently from edges and propagated through interlayer cleavage. Sheet surfaces appear clean and unattacked. (From Giovanoli, Schnoor, Sigg, Stumm and Zobrist, 1-89].)

This may result with certain cations as a conse- quence of oversaturation with respect to surface sites (which occurs whenever the adsorbate surface excess exceeds the density of surface sites or with respect to the solubility of a solid new phase [13,16,33]. For example, in case of Cr(III) the polymeric surface clusters at the surface of hydrous ferric oxide mixed :~- and ,/-CrOOH local structure [16]. Such surface polymerization enhances the cross linkage of the surface lattice and may lead to most efficient inhibition by imparting to the surface a local CrOOH network which is most inert towards dissolution (extremely low water exchange rate).

8.6. Generalizations

The effects observed with iron oxides are similar to those expected to occur with other oxides and

W Stumm/Colloids Surjhces A." Physicochem. En~. ,4,Vwcts 120 (1997) 143 166 163

aluminium silicate minerals; the concepts discussed can be extended to carbonate, phosphate and sulfide minerals.

The organic and inorganic ligands discussed occur all in nature: they can have an effect on the surface controlled dissolution and weathering reac- tions and on the precipitation and crystal growth processes. In natural systems attention must be paid to the competition between dissolution enhancing and dissolution retarding ligands and the blocking of reactive surface sites.

9. Corrosion and inhibition of the dissolution of passive oxide films

Iron. and most other metals, are unstable with respect to their conversion to oxides. Thermodynamically (e.g., by a p e vs. pH diagram) the conditions where corrosion is possible can be elucidated [14]. Solid corrosion products (Fe(OH)2, FeCO 3, Fe(ll , l lI)(hydr)oxides) may form a partially or fully protective coating that retards corrosion: a passive fihn of i ron(II , I l l ) oxide may be formed at high redox potentials (anodic polarization with an externally applied current or in the presence of suitable oxidants such as chromates).

Surface coordination chemistry is essential in interpreting the process occurring in corrosion and its prevention. The solution variables (pH, ligands, redox components) will have a pronounced effect on the type of corrosion products formed. Even in the metallic state inner-sphere surface bonds with ligands can be formed in anodic dissolution [90]. The complex pH and ligand dependence of the dissolution reaction depends on the competitive role of wtrious types of coordination mechanisms such as

(1) H ,O as a ligand:

F'e(s) ~ Fe,~ + e (20)

FGd~-~ Fe(aq )2+ + e

(2) OH as a ligand:

Fe(s) + H 2 0 ---, FeOH~ds + H + + e (21)

FeOH~ds--, FeOH(aq) + + e

(3) Anion A as a ligand:

Fe{s) + A --+ FeAad s q- e

FeAads-+ FeA(aq) + + e (22)

An elaborate interpretation of the role played by FeOHaa S in the catalytic mechanism was pro- posed by Heusler [91]. According to this model, FeOHad s is associated with a kink at a dissolving edge. If one FeOHad s is removed from such a position, the active site for dissolution is simply transferred onto the next atom on the edge.

The subsequent reactions of Fe(aq 12~ , FeOH(aq) + and FeA(aq) + (reactions 20 22) with OH , ligands and oxidants form, upon exceeding the solubility, a variety of corrosion prodncts which, under favorable conditions, may become corrosion protective films. Thin oxide films often are corrosion protective (passivating) when they have the right properties (continuous lilms, with partial electronic conductivity). We may not know exactly the composition of the passive film, bnt it has been suggested that it consists of an oxide of F% .,04 with a spinel structure. The passive layer seems to vary in composition from Fc304 (magne- tite), in oxygen-free solutions, to Fe2.~,-,O 4 in the presence of oxygen. It may also consist of a duplex layer consisting of an inner layer of Fe304 and an outer layer of 7-Fe203. Obviously, the surface of the hydrated passive film on iron displays the coordinative properties of the Fe~lll) surface hydroxyl groups. To what extent is our knowledge on the reactivity of oxides with H ". ligands, and metal ions useful, for the understanding of passiv- ity'? In many cases the dissolution rate of passive metals is related to the dissolution rate of the passive film [70]. Although, the concept of oxide surface reactixity, as influenced by solution wtri- ables does not seem to be a comprehensive part of corrosion and passivation literature, some of the information reviewed here on the structural chem- istry and on the effects of surface complex forma- tion on the dissolution appears to be applicable to the interpretation of some of the factors that enhance or reduce passivity. It is clear that the pH and the nature of the ligands {anions) in solution can determine the growth and development of the surface oxide film. The oxoanions borate, phos-

164 V~ Stumm/Colloids Suffdces A." Physicochem. Eng. A,spects 120 (1997) 143 166

phate, arsenate, selenite, and silicate that have been shown (Figs. 6(b) (d)) to inhibit both non-reduc- tive and reductive dissolution of Fe(III)(hydr)oxides, are known to enhance the passivity of oxide films [76].

For example, the choice of borate as the "ideal" solution for passivating iron is probably due to the good dissolution inhibitive abilities of borate. Anodic activity of iron is found to be much lower in borate solutions than in acetate solution of the same pH (7.4) [92,93]. The oxoanions, in addition to their ability to inhibit dissolution, impart to the iron(hydr)oxide surface a negative surface charge at the neutral pH range [94]. The high electric field established by the potential gradient across the oxide film tends to facilitate the transport of aggressive anions through the film towards the metal surface. The negative surface charge imposed by the surface complexes of oxoanions may retard the migration of C1 and of dissolution promoting ligands and reductants to the Fe surface [95,96]. Oxoanions may also affect the kinetics of oxide film formation, because their presence increases the rate of oxygenation of Fe(II). Oxidizing oxo- anions like CrO 2- and MoO 2 act both as dissolu- tion inhibitors and oxidants maintaining high (passivating) redox potential and preventing reduc- tive dissolution. Any CrO42 that is reduced becomes adsorbed as Cr 3+ or CrOH 2+ [73]. Adsorbed Cr(III) is also an efficient dissolution inhibitor (Fig. 7(b)).

Ligand promoted dissolution of the oxide films assists in the breakdown of passivity and promotes the formation of corrosion pits. Experiments by Strehblow [96,97] with the use of the rotations- ring-disc electrode on the effect of fluoride (0.1 M F - , low pH) show that fluoride promoted Fe(III) oxide dissolution proceeds as follows (using our notation for surface species, cf. Fig. 5)

N / O H N / OH2 Fe Fe

/ \ O / NOH

fast ~ / O H ~ / O H 2 + HF ~ - Fe Fe

/ \ O / \ F + H ~

(23a)

/ O H ~ / O H 2 slow ~ /OH2 Fe Fe + 2 H + ~ Fe

/ \ O / \ F / \ O H + FeF(aq) 2+

(23b)

fast FeF(aq) 2+ + n H F ~ F e F , ( a q ) 3 , + nH +

(23c)

where the detachment step (23b) is rate determin- ing. The same surface controlled mechanism has been proposed by Zuti6 and Stumm [98] for the dissolution rate of the passive hydrous alumina film on an aluminum surface by fluoride on the basis of measurements with a rotating disc alumi- num electrode. Although CI and SO 2- form weaker complexes with Fe(III) than F , a similar mechanism has been proposed for the aggressive action of SO 2- and chloride at low pH [96,99].

10. Concluding remarks

The surface reactivities of the mineral water interface, as reflected in the rate of surface con- trolled processes (dissolution and formation of the solid phase, inhibition of these reactions and cata- lysis) depend on the array of surface complexes formed by the interaction of the surface functional groups with H +, O H - , ligands, and metal ions. Solution based studies on the extent of species binding to the surface have been successful in modeling surface complex formation equilibria and their dependence upon surface charge and deriving proton and ligand dependent dissolution rates. Recent advances made on the appreciation of the fine structure of surface complexes, made above all by X-ray absorption studies, have made it possible to refine our information on the type of linkage of the surface bound species to the reactive sites of the crystallographic surface; thus, it has become possible to distinguish not only between outer- sphere and inner-sphere surface complexes, but also between mono-nuclear and bi- or multi- nuclear surface species and in some cases also to differentiate between varying surface reactivities at different crystallographic planes of the solid sur- face. This detailed structural information is desir- able to deduce the configuration of the precursor of the activated complex for the growth or dissolu- tion reaction and is essential for revealing mecha- nisms of inhibition. It can be shown that bi-nuclear (or polynuclear) surface complexes, especially those formed by oxoanions such as phosphate, arsenate,

W. Stumn(Colloids Surfaces A: Phvsicochem. Eng. Aspects 120 ~ 19971 143 166 165

s i l icate , b o r a t e a n d sulfa te , a n d by m a n y m u l t i -

v a l e n t c a t i o n s s u c h as V O 2+, A1 s+, a n d C r 3+, t h a t

b r i d g e t w o (or m o r e ) d i s s o l u t i o n a c t i v e su r f ace

m e t a l c e n t e r s of t he la t t ice , a re p a r t i c u l a r l y g o o d

a t i n h i b i t i n g r e d u c t i v e a n d n o n - r e d u c t i v e d i s so lu -

t i o n o f i r o n ( I I ) o x i d e s . T h i s p h e n o m e n o n is

be l i eved to be d u e to the l a rge a c t i v a t i o n e n e r g y

w h i c h m u s t be o v e r c o m e to s i m u l t a n e o u s l y r e m o v e

t w o su r face m e t a l cen te r s . U n d e r s t a n d i n g the sur -

face c h e m i s t r y a n d the r e a c t i v i t y a t t he m i n e r a l -

w a t e r i n t e r f a c e i m p r o v e s o u r a p p r e c i a t i o n of ele-

m e n t a n d p o l l u t a n t t r a n s f e r a t t he e a r t h ' s su r f ace

[ 100] . T h e o x i d e d i s s o l u t i o n p r o m o t i n g effects of

H + a n d l i g a n d s a n d t he i n h i b i t i o n effects o b s e r v e d

by bi- o r m u l t i - n u c l e a r su r face c o m p l e x e s a re a l so

of i m p o r t a n c e in t e c h n o l o g i c a l s y s t e m s s u c h as

m e t a l c o r r o s i o n a n d its i n h i b i t i o n .

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