Eawag_06423

Corrosion of Metals in Aquatic Systems; an Introduction

Werner Stumm

EA WAG Swiss Federal Institute for Environmental Science and Technology Ueberlandstrasse 133 CH-8600 Duebendorf (Switzerland) e-mail: stumm@eawag.ch

Corrosion of Metals in Aquatic Systems

Abstract

Metal corrosion is interpreted as an electrochemical process from both a thermo-dynamic and kinetic point of view. Thermodynamics defines the conditions (pc, pH, solution variables) under which corrosion of a metal is possible and whether corrosion products are soluble or insoluble. Thermo-dynamic information is also used to understand the principles of cathodic protection and anodic stimulation ("local batteries"). The elementary aspects of the kinetics of electron transfer at metal electrodes is reviewed and it is illustrated how polarization curves on corroding specimens are measured and how polarization data can be used to measure corrosion rates and to interpret the mechanism of corrosion and its inhibition.

Often the CaC03 formed on the surface of the corroding iron walls, together with some corrosion products, may provide some protection and retard corrosion. The question of CaC03 deposition, however, is not only one of equilibrium (saturation index). The cathodic reaction, the reduction of 0 2 or H+, on a corroding iron surface produces a pH increase in the proximity of the corroding metal. Although the maintenance of a saturation index (SI) near zero or slightly positive has often been expedient to reduce corrosion and to prevent a clogging of the pipe by CaC03, the problem of the mitigation of corrosion and "red water" cannot be solved by this simple recipe alone.

The corrosive behavior of a few metals is essentially determined by the kinetics of the dissolution of the corrosion products. This seems to be the case for Zn in HC03 solutions, for passive iron in acids, and for passive Al in alkaline solutions. The mechanism of the dissolution of iron and of the passivation of this dissolution is extremely complex. The exact composition of the passive film is not known; but it has been suggested that it consists of an oxide of Fe3_x04 with a spinel structure. The passive layer seems to vary in composition from Fe30 4 (magnetite), in oxygen-free solutions, to Fe2.670 4 in the presence of oxygen. The dissolution of passive oxide films is compared with the dissolution of hydrous metal oxides and the essential features of the concept of proton- and ligand-promoted dissolution of hydrous oxides (Stumm, 1997) is used to interpret the characteristics of passive oxides and their dependence on solution variables.

This introduction may serve to direct readers to some of the more detailed literature on the chemistry of corrosion.

1

1.1 Thermodynamic Aspects

Most metals are, with regard to their conversion to oxides, unstable. The reaction xM + y/2 0 2 = MxOy is characterized by the following free energy values:

Oxide AG0 (kJ mo1-1) Oxide AG0 (kJ mol-1)

Fe20 3 -742.3 CuO -129.7 Al203 -1582.4 NiO -211.7 Cr20 3 -1058.2 ZnO -318.4 MgO -569.4 Sn02 -519.7

The electrode potential1> EH (or pE)2> vs pH diagram for the various metals gives the predominance areas where corrosion, from a thermodynamic point of view, is possible. In case of Fe we may con-sult Figure 1. The non-shaded area of predominance of Fe2+, Fe3+ and FeOH2+ is the thermodyna-mic range where corrosion occurs. If the pE of iron can be kept below -10 (EH< -0.6 V), e.g., by applying some external voltage, it will not corrode. In the shaded predominance area of solid pre-cipitates (Fe(OH)i(s), FeC03(s), Fe(OH)3(s)) corrosion is still possible but the precipitate may under suitable conditions form a partially or fully protective coating that retards corrosion; such a passive film of iron(ill) oxide may be formed at high pE (EH) values (anodic polarization or in the presence of suitable oxidants such as chromate).

The electrochemical series (Table 1) gives thermodynamic information on the so-called "nobility" of various metals; the higher the standard electrode potential, the more noble is the metal, silver being more noble than Cu and Cu being more noble than Zn.

Table 1 Electrochemical Series of a few Metals (25 °C)

Reaction

Mg2+ + 2 e- = Mg(s) zn2+ + 2 e- = Zn(s) Fe2+ + 2 e- = Fe(s) 2 H+ + 2 e- = H2(g)

Cu2+ + 2 e- = Cu(s) Ag++ e- = Ag(s)

Standard Electrode Potential (V)a)

-2.35 -0.76 -0.44

0 0.34 0.8

logK

-79.7 -26 -14.9

0 11.4 13.5

a) The standard electrode potential (also called standard redox potential) is the potential that would be observed if all substances in the redox reaction were in their standard state of unit activity; it is related to the free energy change for the cell reaction, .6.G0, or the equilibrium constant of the reduction reaction, E~ = .6.~o = ~ In K.

1 > EH refers to the electrode potential (volt) in reference to the standard hydrogen electrode (see Figure 2). 2> pe defines a hypothetical electron activity: pc =-log {e-}. pe is related to EH by pe = EHF/2.3 RT. F =Faraday

(96,485 Coulombs mor 1) (Stumm and Morgan, 1996). 2

G;. Potential (V) a) +20 1.2 b)

-.,._ ...... ......... Po >1 atm 2 ....... 2 ..... '·,,

+10 0.6 +20

Fe2+ +10 ...................... _.... .... w

0.. 0 0

pE EH .......

....... 0 ..._

p~>1 -10

-10 -0.6 -1

Fe(s) -20

3 6 9 12 0 7 pH pH

Figure 1 a) Diagram pe or E8 (electrode potential vs standard hydrogen electrode) versus pH for the system Fe-C02-H20.

The solid phases are Fe(OHh(s) (amorphous), FeC03(s) (siderite), Fe(OH)i(s), and Fe(s). Lines are calculated for Fe(Il) and Fe(ill) = 10-5 M and CT ([H2CO;'] + [HCO;J + [CO~-]) = 10-3 M (25 °C). The formation of Fe30 4(s) has not been considered; it may be formed in the predominance area between Fe(OH)i(s) and Fe(OH)3(s). The possible conversion of carbonate to methane at low pe values is ignored.

b) A simplified potential-pH diagram (often referred to as Pourbaix diagram) indicating domains of EH potentials where corrosion is thermodynamically possible and impossible (immunity). In the domain where Fe(ill)-(hydr)oxides or other solid phases are stable, partial or full protection from corrosion ("possible passivity") may be possible. For derivation of pc: or EH vs pH diagrams see Stumm and Morgan (1996).

We may compare the following (hypothetical) cells: H2(g) I H+ 11 Cu2+ I Cu(s) Potential difference 0.34 V Fe(s) I Fe2+ 11 H+ I H2(g) Potential difference 0.44 V Mg(s) I Mg2+ 11 Fe2+ I Fe(s) Potential difference 1.91 V Fe(s) I Fe2+ 11 cu2+ I Cu(s) Potential difference 0.78 v

These cells are written in such a way that electrons in the external circuit flow from left to right. Figure 2 exemplifies the principle of an electrochemical cell.

Cathodic Protection and "Local Batteries". If an iron metal is connected (external circuit) to a Mg metal, the latter becomes an anode and Mg2+ goes into solution; the iron becomes the cathode and is thus protected by the electrons delivered to the iron from the Mg-electrode. One speaks of cathodic protection. A similar situation arises with a Zn-coated iron metal. As long as the Zn acts as an anode (dissolves), the iron is cathodically protected. Cathodic protection can also be achieved by applying some voltage between a counter electrode and the iron. Alternatively, the connection of Cu with Fe makes the copper the cathode and the iron the anode; the dissolution (corrosion) of Fe is enhanced. A similar phenomenon occurs if the water flowing through the iron pipe contains Cu(II). According to the electrochemical series, Cu(II) will be deposited on the iron metal surface; the resulting "local battery" causes iron to corrode. In this case often a pitting corrosion is observed.

A corroding metal is like a short-circuited energy producing cell. The half-reactions responsible for the corrosion of iron in water are, for oxidation,

3

and, for reduction, either

or

0 2(g) + 4H+ + 4e- ~ 4H20

(02 + 2 H20 + 4 c --.. 4 OH-)

2 H+ + 2 e- ~ H2(g) (2 H20 + 2 e- --.. H2(g) + 2 OH-)

These half reactions occur each in a sequence of elementary reaction steps; one of these steps is usually rate determining. The Fe(II) formed can become oxidized to Fe(ID). Figure 3 illustrates schematically how anodic and cathodic processes occur simultaneously on a corroding iron surface.

In order to understand the electrochemical mechanisms of corrosion in electrochemical terms and how to express corrosion rates in terms of current density we briefly review some elementary aspects of electron transfer at electrodes.

H2(g) ---.

Figure 2

e- .... negative

nega1i.ve terninal electrode

j~

Ult2{Jt i~1~ill i~: H2(g) = 2 H+ + 2e- ~·.::. · .\· .. ·:

\X/\{//~1-;~~t?: {H+J = 1 M PH2= 1 atm

SHE

E=0.34 V e- ....

positive teminal positive

e-l electrode

Salt bridge

(KO)

Electrochemical cell with a standard hydrogen electrode (SHE) as the left electrode. By convention, the negative terminal of the meter is on the left and the positive terminal is on the right. Hydrogen gas at 1 atrn pressure is being bubbled through the solution near the left electrode (Pt) to keep the solution phase activity of H2 equivalent to 1 attn. {W} in the left cell is 1 M. The rifht electrode is made of Cu(s). { Cu2+}in the right cell is 1 M. At the left electrode, H2 is being oxidized to 2 H . At the right electrode, Cu2+ is being reduced to Cu(s). Arrows indicate the direction of the flow of electrons. The voltmeter (potentiometer) measures a positive voltage of 0.34 V. (The measurement has to be made when the cell does no work (net current i = 0)) (Adapted from Pankow, 1991)

4

Figure 3

Anode Fe-..... Fe2+ + 2 e-

Scheme of anodic and cathodic reactions on corroding iron surface in presence of 0 2• The iron metal is the conductor of electrons between local anodes and cathodes; the electrolyte is the ionic conductor. In cathodic areas, the cathodic reduction of O, consumes Ir (or produces OH-). Of course, the rate of electron production equals the rate of electron consumption.

1.2 Electrode Kinetics

At an electrode-electrolyte interface a charge separation between the metal surf ace and the electrolyte occurs. Let us first consider the transfer of an electron between Fe3+ and Fe2+ in an acidified solution with the help of an inert (e.g., Pt) electrode.

To cause the passage of a finite current at this electrode it is necessary to shift the potential from its equilibrium value, by applying a small voltage between the auxiliary electrode and the iron electrode. We thus obtain a curve depicting the electrode potential as a function of the (applied) current (polarization curve). At the equilibrium potential, that is, at the point of zero (applied) current, the half reaction

is at equilibrium, but the two opposing processes, the reduction of Fe3+ and the oxidation of Fe2+, proceed at an equal and finite rate,

v 1 (rate of reduction) = v2 (rate of oxidation) (1)

which is proportional to the rate of passage of electrons in both directions. Although the net rate of passage of electrons is equal in both directions and thus the net (applied) current is zero, the passage of current in a single direction is not zero and is called the exchange current, i0 (see Figure 4).

The rate of electron exchange at the surface of the electrode may be expressed as moles reacting per unit time per unit surface area. Under simplifying assumptions (first-order kinetics) v1 or v 2 will be proportional to the concentration or activity of the reactants at the surface of the electrode ([Fe3+]e and [Fe2+le)· (For other systems we may substitute [ox]e and [red]e, respectively.)

V1 = ki [Fe3+]e v2 = k2[Fe2+]e

(2a)

(2b)

where k1 and k2 are velocity constants (with the units cm·1 sec-1) of reduction oxidation and oxidation, respectively, and where concentrations are conveniently expressed as moles cm-3. Velocity constants are a function of activation energy, Eact,

k = Z exp (-Eact/R T) (3)

5

In the case of electrochemical reactions, which are rate-controlled by electron transfer, the activation energy is a function of the difference EH - ~ between the potential of the electrode, EH, and the standard oxidation potential, E~, of the redox system (hydrogen scale) in question. It can be shown that

(crnF(EH - E~))

ki = ko exp - RT (4)

and

(5)

where n is the number of electrons required in the kinetic process (n = 1 for the Fe3+_pe2+system)lJ and a is called the transfer coefficient and is a measure of the symmetry of the energy barrier; a is the fraction of the total energy that acts to decrease the height of the energy barrier for the reduction reaction and 1 - a is the fraction that tends to increase the height of the anodic potential barrier. (a is frequently close to 0.5.) Fis the Faraday (96.485 C mo1·1 ).

Because current is, proportional to the velocity of reduction or the velocity of oxidation, it follows that

ic = nFv1A ia = nFv2A

where A is the area of the electrode. The cathodic and anodic currents can then be given by

(6) (7)

(8)

(9)

The behavior of each of these currents as a function of the cell emf is shown in Figure 4. The net current, i, is the sum of the two opposing currents ia and ic. Adopting the convention ia > 0 and ic < 0, i is given by

(10)

Thus, at any potential (these may now be values measured with an electrochemical cell, i.e., not equilibrium values) the net current is given by

(11)

This type of electrode kinetic equation is often referred to as the Butler-Volmer Equation. At equilibrium where no net current flows (i = 0, ia = ic), concentrations (activities) at the electrode

1 > Multielectron transfer reactions proceed often in subsequent one-electron transfer steps. 6

surface equal bulk concentrations and potentials measured correspond to equilibrium conditions, ( 10) and ( 11) yield:

(12)

or in a more general sense

(13)

Eqs (12) and (13) are identical with the Nemst equation.

The exchange current i0 is given by ic or ia at equilibrium. By combining (8) or (9) and (12) the exchange current i0 becomes

(14)

and (11) can be rewritten as:

i = io(e-<anF/RT)liE _ e(l-a)(nF/RT)LIB) (15)

where LIB is the deviation in EH from its equilibrium value that results from the passage of a net current. ~ is often referred to as overvoltage. When .6.E is very small, the exponential terms in ( 15) can be expanded in the form e-x::::: 1 - x and (15) becomes

i = -i0(nF/RT)Llli (16)

Corresponding! y the slope of the polarization curve, (Ail LIB )i = 0 is given by

(17a)

The equations given above are illustrated in Figure 4. Eq ( 15) is often referred to as the Butler-Volmer equation which describes the general relationship between current density and potential for a simple electrode reaction controlled by charge (electron) transfer.

7

[Fe3+] = 10·3 M [Fe2+] = 10·3 M

pH=2

a)

I ,. I I

I I I I

I I I I

I • I I 10 I

i(µA)

100

, I Fe2+-Fe3+ ,, 1

t1>'"' I

50 +

"""' ... -,.,. ... ...... ..,., I

475 : I/ I . : . ,,, t-'o ,'~50 I I I I I I I I; I , F 3+ F 2 t/ e _.,.. e +

/ ilOO /'' net current ' i(j.tA)

net current

b)

[Fe'+] = I0-4 M [Fe2+] = 10-4 M

pH=2

Fe2+-pe3+ ,,,.,,.""' ---, ,

i(µA)

100

50 +

50

i(µA)

Fe2+-Fe3+ ,,,. .,.,,.

c)

,, , ,,

.,., ,, _.,.

42.f "" .. f Aso 475

[Fe3+] = 10"4 M [Fe2+] = Ifr3 M

pH 2

500

,-;"""" ,'

I I • •.... ----·····J __ .....• net current

,/ Fe3+-Fe2+

i(µA)

50 +

50

i(µA)

Figure 4 Polarization curves for various concentrations of Fe2+ and Fe3+. Solid lines =polarization curves (electrode area= 1 cm2); i0 = exchange current. Dashed lines are hypothetical cathodic (-i) and anodic ( +i) currents. Curves are schematic but based on experimental data at significant points. Potential scale is relative to the saturated calomel reference electrode.

8

As indicated, the net current can be visualized as the summation of two opposing currents (cathodic and anodic). The rate of Fe3+ reduction (conventionally expressed as cathodic current) generally increases exponentially with more negative electrode potential values and is, furthermore, a function of the concentration of [Fe3+] and of the effective electrode area. Similar considerations apply to the rate of Fe2+ oxidation (anodic current), which is proportional to [Fe2+], electrode area and the exponential of the potential. It is obvious from the schematic representation that in the case of Fe2+ -Fe3+ an infinitesimal shift of the electrode potential from its equilibrium value will make the half reaction proceed in either of the two opposing directions, provided the concentration of these ions is sufficiently large. The measurement of the equilibrium electrode potential is only feasible in such a case.

Fe2+-.Fe3+ /

I I

I +i

Ecq / ~~~~~t_..,,._ __ +--::::-~----+-...E(V) ----------+

Fe3+_...pe2+

-i

Figure 5 Electrode polarization curves for oxygen-containing solutions with a mixed potential resulting from oxygen reduction accompanied by Fe2+ oxidation. Curves are schematic but in accord with available data at significant points.

Mixed Potentials; Mixed Electrodes The balancing of anodic and cathodic current need not correspond to the same redox process and may be a composite of two or more processes. An example of this is shown in Figure 5 for a Fe3+ -Fe2+ system in the presence of a trace of dissolved oxygen. The measured zero-current potential is that value where the rate of 0 2 reduction at the electrode surface is equal to the rate of Fe2+ oxidation rather than the value of Eeq since at the latter point simultaneous 0 2 reduction produces excess cathodic current. In addition, because the net reaction at EM converts Fe2+ to Fe3+, the measured potential exhibits a slow drift. Such "mixed" potentials are of little worth in determining equilibrium E8 values. Such a mixed potential is often observed if a metal is in contact with an electrolyte where two or more electrode reactions, which are not dependent on each other, talce place simultaneously.

1. 3 Corroding Systems, Polarization Curves

Mixed (non-equilibrium) potentials occur also in corroding systems where the dissolution (oxidation) of the metal

9

Me __. Me+zn + ne-

is balanced by the reduction of water, H+ or 0 2

or

2 H+ + 2 e- -+ H2(g)

(2 H20 + 2 e- -... H2(g) + 2 OH-)

0 2 + 4 e- + 4 H+ --.. 2 H20

(02 + 4e- + 2H20 ---.... 40H-)

The cathodic (reductive) and anodic (oxidative) current reflect rate detennining reactions at the corroding surface (Figures 6a and 6b).

a) [Acm-2

ianodic

Mel++ 2 e- -Me

icathodic

b) i [Acm-2]

anodic

cathodic

Me --Me2++ 2 e-

2 JI++ 2 e- -H2(g)

2 H20 + 2 e- -2 OH-+ H2(g)

E(V)

Figure 6 a) Schematic representation of the dissolution of a metal by H+. Oxidation of metal is accompanied by reduction

of H+. Mixed potential occurs where anodic current density = cathodic current density = icorr = corrosion current density. The experimentally measured net current density, inet' is the sum of the partial anodic and cathodic current densities 1.

b) Simplified scheme of pH dependence of corrosion rate of an active metal by H+ (H2-formation). In reality the situation is often more complicated. Ligands may influence the anodic dissolution reaction, e.g., at high pH hydroxo complexes Zn(OH)_r, Pb(OH)i- may be formed.

Noble and Active Metals. Figure 7 displays schematically the corrosion of noble and active metals in electrolytes containing dissolved oxygen. Oxidants other than oxygen also facilitate corrosion. Increase in the concentration of the oxidant tends to increase the corrosion rate (unless insoluble oxides are formed). Complex formation of ligands with the dissolving metal ion shifts the current curve of the dissolution to more negative potentials and thus, enhances usually the corrosion rate. In alloys the binding of the dissolving metal leads to a positive shift of the current curve of dissolution. It should be noted that Figures 6 and 7 are all schematic, illustrating in a qualitative way the principles involved. The phenomena of passivity which will be discussed later are not yet considered.

JO

i [A cm-2]

anodic

cathodic

Figure 7

(Zn, Fe, Pb)

I I I

smaller 1 I I

(Cu)

I [02,] I / . - - - - - - - _t_ _ - .. 0 2-Reductton larger

(Ag, Hg. Au)

E (V)

Schematic representation of corrosion; oxidation of noble, less noble and active metals are compared with reduction of 0 2• (At lower potentials the reduction of 0 2 is diffusion limited.) The anodic curves can be influenced by ligands in solution, and by the formation of insoluble oxides or other precipitates).

Tafel Equations. In corroding systems, a modified Butler-Volmer equation (cf. Eq 15) is used; the detailed mechanisms of the partial electrode reactions are frequently not known and one often uses the equation

(anF11) ( (l-a)nF11) = io exp RT - io exp - RT (18a)

. . (11 J . ( 11 J 1 = lo exp Pa - lo exp - Pc (18b)

where 11 = the polarization ( overvoltage). 11 indicates the deviation of the electrode potential from the corrosion potential; where Pa and Pc are defined by Eqs 19a and 20a.

Pa = d E/d ln I ia I and Pc = -d E/d ln I ic I

In a plot of ln I i I vs E we can describe the relationship for large anodic polarization, 11, as

(19a)

where aa = -Pa ln io

and large cathodic polarization as

(20a)

where ac = Pc ln io

Polarization data are often plotted as 11 vs ( decadic) log I i I . Eqs 19 and 20 are then rewritten as

11

a) b) [A crri2J I

I lial I bglil l\;I

anodic / i i a a/ I ba I

/ I I

/ tcxr I

/ I I .,,.., I

} , , .. E '-.Ji ------ logi"'rr I

C«T ,' I , , I , I .; , I

; I ..... , I ; I

' I catha:iic ; I

I

~: ic /ic ' E , , ,

.I I

I I

Figure 8 Scheme of arithmetic (a) and logarithmic (b) presentation of current potential curve of a mixed electrode showing the corrosion potential, Ecorr• and the corrosion current density, icorr· The drawn-out line is the polarization curve (measured net current density vs electrode potential). The broken lines indicate the partial current densities, anodic for the oxidation reaction (the dissolution of the metal), and cathodic for the reduction reaction (reduction of H+, H20 or oxygen); the value of icon can be obtained by extrapolating the anodic or cathodic partial current density to the corrosion potential. ba and be are the Tafel slopes for the anodic and cathodic current densities.

Figure 9

applied vo1tage

~11---1 --~_,,..._,,..,._,,..,._~~-,;l"·I

reference ---.. !®--electrode

capillary

test electrode

auxiliary electrode

Principle of experimental arrangement for measuring the polarization curve of a (corroding) metal specimen. An imposed potential difference drives the current from the auxiliary electrode to the test electrode. Usually a potentiostatic circuit is used. The electrode potential is measured with a reference electrode which is brought close to the test electrode (with the help of a capillary salt bridge) to minimize iR drop (ohmic resistance) in test solution.

12

(19b)

where a~= -2.303 Pa In i0

and

(20b)

where a~ = 2.303 Pc In i0

The so-called Tafel slopes ba and be (from a plot T\ vs log I i I) are related to Pa and Pc (slope T\ vs In I i I) by

ba = 2.303 Pa be = 2.303 Pc

(21a) (2lb)

The polarization curve (Figure 8) is obtained experimentally if a voltage is applied between the corro-ding metal (used as test electrode) and an auxiliary electrode, and the net current density is measured as a function of the electrode potential (Figure 9). The partial anodic and cathodic current densities as a function of the electrode potential are inferred from the mechanism of the anodic and cathodic process.

I. 4 Evaluation of Polarization Curves Polarization Resistance inversely related to Corrosion Rate

In Eq ( 17) we derived for the simple case of electrodic electron transfer in the Fe3+ + e- = Fe2+ system that the exchange current, i0, is related to the slope of the polarization curve at zero overvoltage.

(17b)

The slope of the polarization curve is the reciprocal of the polarization resistance rE = (6:) [ohm cm2]. 1 = o

In a mixed electrode or corroding system the current at zero overvoltage, i.e., at the corrosion potential, EcorP is the corrosion current density Cicorr). The corrosion current density is inversely related the polarization resistance, i.e., proportional to the slope of the polarization curve Cinet vs EH). The slope (Lii/ ~)i = 0 can relatively easily be measured by measuring the current response as a con-sequence of small shifts of the electrode potential (a few mvolts). The measurement of the slope of the polarization curve can thus be used as a measure of the corrosion rate (with relatively little disturbance of the corroding specimen). Figure 10 illustrates the relationship between the slope of the polarization curve and the corrosion current.

In ideal cases the following relationship can be derived:

13

(22)

where ba and be are the slopes of the ( decadic) logarithmic partial anodic and cathodic polarization curves (Tafel slopes), respectively.

Eq (22) shows that the corrosion current density is determined by the kinetic parameters of the partial electrode reactions. Eq (22) is the basis for the empirical relationship often observed between corrosion rate and the reciprocal of the polarization resistance.

. (LliJ -1 lcorr = f AE i = 0 = fr E (23)

where f is a proportionality factor, and rE is the polarization resistance.

Figure 10 gives an experimentally determined correlation between instantaneous corrosion rates (as detennined from weight loss measurements of cast iron specimens exposed to different natural waters for ~it£'erent time periods) and the reciprocal of the polarization resistance, (di/ LlE)i = 0. The factor µb

= b a+ ~ is of such a form that in a given system, µb remains relatively unchanged for slight changes a c

in individual b values. It is thus explainable that fin Eq (23) is a rather constant proportionality factor that can be determined empirically in a given corrosion system. In a corroding system where the corrosion rate is controlled by a limiting cathodic current (e.g., transfer of oxygen to the cathodic

b areas) the factor fin Eq (23) becomes~ (be in Eq (22) is set equal to infinite).

It is essential that the measurement of polarization resistance be measured in such a way that the measurement does not include an ohmic resistance

where rcorr is the polarization resistance due to corrosion and r0 = d.6. <llofdi, the ohmic resistance. By placing the reference electrode in close proximity to the corroding electrode (usually with the help of capillary salt bridge; see Figure 9) the ohmic resistance in the test solution can be minimized. Special electronic measuring techniques are available to make measurements that minimize the ohmic resistance.

14

40

10

1 icorr~~i=0(0.034) Acm-2

2 4 6 8 10 12 14

(~\_ x IQ4 ohm-1 cm-2 \~1=0

Figure 10 Corrosion rate, as measured from weight loss of cast iron specimens exposed to different natural waters, is directly related to the reciprocal of polarization resistance: icorr = (Lii/.1E)i = 0 x 0.034 A cm-2. (From Stumm, 1959)

Polarization Curve as an Aid to Understand Corrosion and its Inhibition. Many of the factors affecting the rate of corrosion can be understood from a graphical superposition of the anodic (metal dissolution) and cathodic (reduction) reactions. The polarization curve can be obtained either experimentally or from a know ledge of the parameters that determine the kinetics of the anodic and cathodic reactions and their inhibition. The semi-logarithmic diagram (log Ii I vs E) is often referred to as an Evans diagram. Sketches of such diagrams are given for various ways of rate control for the overall corrosion process (Figure 11). Schematic Evans diagrams assist in identifying the type of control over the corrosion current and the possibility to influence the corrosion mechanism by applying suitable inhibitors. As Bockris and Reddy (1970) point out, the addition of P, As, or Sb, e.g., as As20 3, reduces the charge transfer for hydrogen evolution (cathodic partial process) and thus reduces corrosion rates (Figure 11 b ). Or if oxygen is the reaction at the electron source (cathodic) areas, then by adding substances which react with dissolved oxygen, the 0 2 concentration is reduced and thus, the current due to the cathodic current is reduced. The same effect is obtained by films or solid products of the metal dissolution (Fe(ill)(hydr)oxides, FeC03) which retard the transport of 0 2 to the cathodic areas (Figure llc). In natural water systems HCOj ion acts as an inhibitor because it interacts with the OH- ions formed in 0 2-reduction and with Ca2+ to form CaC03 over the electron source areas (CaCOrFe(ill)(hydr)oxides) coating in water supply systems. Furthermore, a film formed may produce a general blockage of the metal surface and may reduce the cathode area fraction or the anode area fraction or both.

15

a) log Ii I

b) log! i I

c) log Ii I

t I:: t 0 0 0

:: (.)_ 0 u J:.Il J:.Il P-1

E(V)

E (V)

Figure 11 Simplified log I i I vs E diagrams to illustrate the role of rate control and influence of inhibitors. a) reduction of corrosion rate by decrease of anodic reaction rate and shift of Ecorr to higher values; b) reduction of corrosion rate by decrease of cathodic reaction rate and shift ofEcorr to lower values; c) reduction of corrosion rate by limiting transport, given by iL, of reductant (e.g., dissolved oxygen) to the

su..-face (icorr S: iL).

However, many organic or inorganic coatings are highly permeable for water and oxygen; thus, it is often not the barrier effect on the diffusion process that gives rise to the corrosion inhibition, but the specific electrochemical properties of the adsorbate metal interface; many adsorbed inhibitors reduce both anodic and cathodic partial reactions of the corrosion process. An efficient organic inhibitor should be able to bind to the metal surface, e.g., the reaction of silanes with hydroxylated metal surfaces (Stratmann et al, 1995).

H 0

H H 0 0

I I I +RSiX3'H20 --· -HX

I I I [-SiO SiO Si-]

J ~ d I I I

16

Alternatively compounds which sorb on the electron-sink areas of the corroding metal slow down the metal dissolution reaction (Figure I la). The compounds most often used for this purpose are nitrogen-containing organic compounds (amines) and sulfur-containing compounds (thiurea) and various aldehydes. (Regrettably most of these substances are not suitable for use in water supply systems.)

The corrosion rate and the corrosion potential are the measurable quantities in the steady state of spontaneous corrosion. The extent of the variation of corrosion potential with time and with corrosion rate is indicative of the mode of inhibitor action, e.g., when reduction of corrosion rate with time is accompanied by a decrease in corrosion potential, cathodic inhibition is more prevalent than anodic inhibition.

Fe-Corrosion. 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 (Hensler, 1982). The complex pH and ligand dependence of the iron dissolution reaction depends on the competitive role of various types of coordination mechanisms such as

I) H20 as a ligand: Fe(s) -... Fe~ds + C

Fe~ds -... Fe( aq)2+ + e-

2) orr as a ligand: Fe(s) + H20 -... FeOHads + H+ + e-FeOHads-... FeOH(aq) + + e-

3) anion A- as a ligand: Fe(s) + A- -... FeAads + e-FeAacts -... FeA( aq) + + e-

(25)

(26)

(27)

An elaborate interpretation of the role played by FeOHads in the catalytic mechanism was proposed by Heusler (1982). According to this model, FeOHacts is associated with a kink at a dissolving edge. If one FeOHads 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)2+, FeOH(aqt and FeA(aqt (reactions (25) - (27)) with OH-, li-gands and oxidants form, upon exceeding the solubility, a variety of corrosion products which, under favorable conditions, may become corrosion protective films. Thin oxide films often are corrosion protective (passivating) when they have the right properties (continuous films, with partial electronic conductivity).

An example of partial inhibition of iron corrosion in oxygen containing solutions is given in Figure 12 (Morsi et al., 1980). The current potential curves given were obtained on rotating iron specimen disks in presence of oxygen (air saturation) at pH 8. This example should briefly illustrate how a corrosion investigator may proceed in order to gain from polarization data insight into inhibitory processes. Figure 12a shows how corrosion rates decrease with time; the decrease is accompanied by a decrease in corrosion potential, thus signaling primarily a cathodic inhibition caused by corrosion products

17

which impede the transfer of oxygen to the cathodic area. Note the line for the limiting current for oxygen supply to the surface. In Figure 12b the presence of 10-2 M HPOl- accentuates this inhibition; the presence of HPoi- causes a more pronounced inhibition than the Fe(ill)\hy?r)o~de products alone. In Figure 12c the effect of a commercial inhibitor on a phosphate basis 1s displayed. The reduction of corrosion rate is now accompanied by a shift in corrosion potential to less negative values indicating rate control by primarily anodic inhibition. In this case pitting corrosion is observed. The pitting could, according to the authors, be avoided in presence of sufficient alkalinity (HCOj) in the corroding solution.

The oxygen that is reduced in a corroding iron system is not solely reduced on the non-rust covered "cathodic" iron surfaces but is also reduced by the Fe(II) incorporated into the oxide layer. This implies that the electronic structure of the oxides will strongly influence the reduction of oxygen and therefore also the corrosion rate (Stratmann and Muller, 1994).

1. 5 The CaC03-Fe(Il,III)(hydr)oxide Coating in Water Supply Distribution Systems

Fully passive oxide films usually cannot be formed in natural waters. But often the CaC03 (calcite) formed on the surface of the corro.ding iron walls, together with some corrosion products, may provide some protection and retard corrosion. Equilibria for the CaCOTsolubility

CaC03(s) + H+ = Ca2+ + HC03 CaC03(s) + H2C03 = Ca2+ + 2 HCOj

can describe under what conditions CaC03 is deposited (saturation index1J). Equilibrium constants for *Ks are given in Table 2. The problem, however, is more complicated because the question of CaC03 deposition is not only one of equilibrium (saturation index). The cathodic reaction, the reduction of 0 2 or H+, on a corroding iron surface produces a pH increase in the proximity of the corroding metal. Although the maintenance of a saturation index (SI) near zero or slightly positive has often been expedient to reduce corrosion and to prevent a clogging of the pipe by CaC03, the problem of the mitigation of corrosion and "red water" often cannot be solved by this simple recipe alone.

It is generally more difficult to protect corrosion in a water of low alkalinity (HCOj) and low [Ca2+] because of the low buffer capacity of such waters (Stumm, 1960); even at a SI z O such waters produce large pH differences at the cathodic and anodic surfaces of the corroding metal.

l) CaCOrsaturati~n is oft~n. as.sessed. by a saturation index (SI = pHactual - pH5) by comparing the actual pH with that a~ hypothetical eqmhbnum with CaC03. If the saturation index is positive (pHactual > pHs) CaC03 will deposit.

18

a)

t ~ a u x

<c: ....._

b)

t N s u x

<c: ....._

c)

t ~ s u x <c: ....._

Figure 12

5 I I

2 I I --+.-----.+-----, I I

10-3 ' J • I _::=_-::~Id, Oz-f------

5

2 10-4 5

2

2 I I I I I I I • I "'

-- _-:::- -i =--=--::-:::--= --= =-1~ =1ct, 0:0: ~-"' l : "'~/ --- I ~ I -------~-- ------:--- ------:.+- ,~:-__ ~

I I I .

10-3

5 : : : 1corr 1

I I I I I

-------~-------,-------- ----2 1 I I I I I

I I ------ _.!.._ ------·--------~ -

I I I : I L' • I I /I I

-------~--------, ""::---7- -------~-------' I , , "'1 I

: : ...,_ :2h : I I I I

10-4

5

I I I I -------~-------~--- -- -r-------,-------2

wi th 10-2 M HP02-' 4 10-5 ..__ __ __.._ __ ___._E --l-E

-800 -700 -600 cor:._500 cor~OO -300

10-4

5

2 10-5

5 2

10-6

5

-----,-I I

-650 -600 -550

EH/mV--..

Current density-potential curve for iron in the vicinity of the corrosion potential. pH = 8, electrolyte: 0.5 M Na2S04; 1.76 x 10-4 M 0 2 (aerated). a) decrease in corrosion rate accompanied by decrease in corrosion potential; b) presence of 10-2 M HPo;- inhibits corrosion; decrease of corrosion potential; c) presence of a commercial phosphate containing inhibitor "Aktiphos": the decrease of the corrosion rate is

accompanied by an increase in corrosion potential. (Adapted from Morsi et al., 1980)

19

Table 2 Equilibrium Constants for the ReactionaJ CaC03 (calcite) (s) + H+~ HCOj + Ca2+ ~ *K5

5 20 25 40 log *K5 (I= 0) 2.20 2.13 2.06 1.99 1.91 1.69

aJ The following approximation can be used to correct the equilibrium constant for activity (I= ionic strenghth): ' . Vi log ¥Ks (I) = logli< Ks (I = 0) + 2.5 _ r:.

where * K5 (I) [Ca2+J [Hco;1

{ft}

1 + -VI

Example: Adjustment of a ·water into CaC03 saturation equilibrium for corrosion and scale control A groundwater (10 °C) has the following composition: pH= 7.2, HCOj = 3.1x10-3 M, Ca2+ = 1.4 x 10-3 M, I = 5 x 10-3 M.

How should this water be treated to attain CaC03 solubility equilibrium?

CaC03 (calcite) solubility can be characterizeed by the equilibrium

CaC03(s) + H+ ~ Ca2+ + HCOj ; log *Ks = 2.3

The constant given is valid for 10 °C and I= 5 x 10-3 M.

[Ca2+] [HCOj] log {Hi} = 2.3

{ H;} is the hypothetical W ion activity which corresponds to the equilibrium with the actual concentrations of [Ca2+] and [HCOj].

-log {H;} = 2.3 -· log [Ca2+] - log [HCOj] pH5 = 2.3 + 2.85 + 2.51 = 7 .66 Saturation index SI = pHact - pH5 = 7.2 - 7.66 = -0.46.

The water is aggressive (it dissolves CaC03). The adjustment can be made by addition of a base, e.g. NaOH, Na2C03 or Ca(OH)i.

After addition of 3 x 10-3 M Ca(OHh the composition of the water shifts to [Ca2+] = 1.7 x 10-3 M, [HC03] = 3.7 x 10-3 M, pH= 7.5.

The new calculation gives

pH5 = 2.3 + 2.77 + 2.43 = 7.5 SI = 7.5 -7.5 =0

20

"Red" Water

If the products of iron corrosion do not remain attached to the pipe walls, they are released into the water supply; iron release can lead to the appearance of rusty or red water at the tap. The iron release rate may not simply be related to the instantaneous corrosion rate, but depends on a combination of physical and water quality factors. Red water occurs often in distribution systems that are subject to cyclic periods of stagnation and flow. Often under conditions of stagnant water, dissolved oxygen concentrations adjacent to the pipe decreases to zero; ferric oxide in the scale is reduced to dissolved ferrous iron. When the flow of the oxygenated water in the pipe resumes, the dissolved Fenn is taken up and oxygenated to ferric(hydr)oxide and forms red water. For coUoid-chemical reasons. dispersed ferric(hydr)oxide (rather than attached to the scale) is more likely to occur in low calcium and low alkalinity waters.

1. 6 Passivation; Passive Oxide Films

Figures 13a and b give a polarization curve for an iron metal which, at high electrode potentials, displays passivity; the anodic dissolution current shows a transition to a "passive" region where the metal is covered by a thin continuous oxide film. The same type of passivity may also occur if the electrode potential is raised by exposing the corrodible metal to a solution containing suitable oxidants (e.g., chromate). Similar passivity phenomena are observed with other metals, e.g., Cr, Ni, Co. Steels alloyed with elements such as Cr and Mo are more readily amenable to passivation {''stainless steel").

The corrosive behavior of some passive metals is essentially determined by the kinetics of the dissolution of the passive oxide film, i.e., of the corrosion products. This seems to be the case for Zn in HCOj solutions, for passive iron in acids, and for passive Al in alkaline solutions. The mechanism of the dissolution of iron and of the passivation of this dissolution is extremely complex. We may not know exactly the composition of the passive film; but it has been suggested that it consists of an oxide of Fe3_x04 with a spine! structure. The passive layer seems to vary in composition from Fe30 4 (magnetite), in oxygen-free solutions, to Fe2.670 4 in the presence of oxygen. Figure 14 represents a schematic model of the hydrated passive film on iron as proposed by Bockris and collaborators (Pou et al., 1984). Obviously, the hydrated passive film on iron displays the coordinative properties of the surface hydroxyl groups.

21

a)

~ 'E u s

0

-2

·~ -4 .s

-6

__ ... .... \ .-' I

,.- I , I

, I ,' I

' I ~I I I I I I

Fe0 ___.... Fe2+ + 2 e-active

,I

Fe in 0.5 M 8iSO 4

~assive 7µA cm 2

2 8i0 _.. 02 + 4 W + 4 e-

ri~~~~~~~~~~~~~~~~-

1:~ hydrogenevolution: :~ possble : 'I .. • ...... • ...... •

,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,.,~~~~~~~~~~~~~~~:

~~~~~~~~~~~~~~'''''''''b'i'~~~~~~~~~~~~~~~~;. oxygen evolution~: ''''''''''''''water sta e,,,,,,,,,,,,,,,... 'bl r

mmmm~~..:..:..:..:..:..:..:..:..:..:..:mm~~m~~~~~·: ... -~~s.s~. ~ .... ~: -0.5 0 0.5 1.0 1.5 2.0

&! (V)

b) 40

20

~

's 0 0 ~

-20

-40

I F<!'-F<i''

2 H+--H2(g)

Formation of passive oxide film

/

Reduction of passive film Fe3+ I Fe2+ - Fe'>

iiiiiiiiiiiiiiiiii'''''''''''''''''''''''''''''''''''''''''''''''' 111111111111111111 ,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,.

~~~~~~~active ~~~~~ mmmmmm~: passive ~~~~m~~mmm~: 111111,,,,,,,11111 ,,,,,,,,,,,,,,,,,,,.,,,,,,,,,,,,,,,,,,,,,,,,,,,. 111.l'llllllllllllll ,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,,. ............................................................................................................... r

-1.0 --0.5 0.0 0.5 1.0 ~ (V)

Figure 13 a) Potentiostatic anodic polarization curve of iron in 0.5 M H 2S04 (based on data by Franck and Weil (1952). As the potenial rises, the anodic dissolution current first

increases following a Tafel relation and then reaches a diffusion-limited current (depending on hydrodynamic flow conditions). Further rise of the potential leads to a sudden decrease of the current to a value orders of magnitude lower when a very thin oxide film is formed on the surface. In the passive region, the presence of the "passive" oxide reduces the anodic current (passive current density) to a small value reflecting a small corrosion rate. In a first approximation often a steady state between the corrosion rate and the dissolution rate may be assumed; thus, the anodic current in the passive region also reflects the dissolution rate of the oxide film.

b) Typical polarization curve (obtained by cyclic voltammetry) in deaerated pH = 8.4 borate buffer (scan rate 5 mV s·1) based on data by Sinniger (1992). In cyclic voltammetry, the electrode potential is shifted (in a given electrode potential range), first anodically and then cathodically with a given speed of potential change.

22

Figure 14

Iron metal electrode Passive film

Schematic representation of the hydrated passive film on iron. (From Pou et al., 1984.

Passivation by Chemical Oxidants. The occurrence of passivity depends on the attainement of a sufficiently positive electrode potential independent whether this is achieved by applying a voltage or by the presence of suitable oxidants in the solution, e.g., chromate, Ce +4, Fe3+, 0 2 (in slightly alkali-ne solutions). The anodic part of the polarization curve is complemented by the partial cathodic curve for the reduction of the oxidant. Typically different cases can be distinguished (Figures 15a and b).

a) b)

Ecorr

E E

Figure 15 Schematic partial anodic and cathodic polarization curves (adapted from Landolt, 1993). a) Distinction between (1) influence of a "weak" oxidant, Ecorr < EP, (passivation potential) reflecting high

corrosion rate, and (2) the influence of a "strong" oxidant, Ecorr > E , reflecting low corrosion rate. b) A case where two conditions for ia = ic prevail, reflecting an unstable passivity. The corrosion potential can

establish itself in the active or passive domain.

23

1. 7 Dissolution of Passive Oxides and Inhibition of Dissolution

To what extent can the dissolution of passive oxide films be compared with the dissolution of hydrous metal oxides? In order to discuss this question, we first summarize the essential features of the concept of proton- and ligand-promoted dissolution of hydrous oxides (Furrer and Stumm, 1986; Stumm, 1992, 1997, and Blesa et al., 1994).

Figure 16 lists some of the factors that promote and inhibit dissolution of hydrous oxides. Very small concentrations of inhibitors can often be effective because it may suffice to block the functional groups of the solution active sites such as the ledge sites or the kink sites. Often competitive equilibria exist between the surface and solutes which block surface sites and, on the other hand, solutes which promote dissolution, e.g., H+, ligands (Biber et al., 1994).

Rate Laws for Proton- and Ligand-promoted Dissolution. The general proton-promoted dissolution rate (e.g., mol m-2 h-1 may be written as follows (Furrer and Stumm, 1986)

(28)

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 detennined by alkalimetric titration. The number n varies between 1 and 3, and often corresponds to the number of protonation steps before detachment, 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 + OH-~ S-0- + H20), the corresponding rate law is given by

(29)

where C~8 is the surface concentration (mol m-2) of deprotonated surface sites (equivalent to bound hydroxo groups) or S-0- groups.

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 whbse functional groups contain two or more donor atoms (e.g., di-carboxylic acids, hydroxy carboxylates, diphenols, EDTA, NTA) and which can form bi- or multidentate mononuclear surface chelates. At high concentrations, monodentate ligands like F-, so~-, and er may enhance the dissolution. 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 concentration of the ligand at the surface, q (mol m-2) (Furrer and Stumm, 1986).

(30)

where kL is the reaction constant.

There is a special type of ligand-promoted dissolution known as reductive dissolution. This case involves the transfer of electrons from the environment, through a reducing ligand or reductant, to an oxide surface, followed by dissolution. The mechanism of reductive dissolution has been reviewed by Hering and Stumm (1990).

24

a) OH O~

"-t.i 'M/ /""- /""-0 OH

/ ~ Enhanceirem of dissolutio'n Inhibition of dissolution

Ofl! OH2 '\. / '\. /

Me Me

/ " / " OW OH~

"'- /o\ /o-J-Me Me

/ "o/ "oH Surface protonation

Surface deprotonation

Figure 16

OH 0 '\. / '\. /) Me Me -/ "'-o/ "a

Surface comP,ex for-mation witli Ii gand

that form mooonuclear surface complexes, e.g~

oxalate, salicY.late, citrate, d1phenols, etc.

OH OH2 '\. / ' / Fe11l Fell / "-c{ "oH

Redoced lattice sur-face ion (e.g., Fe(II)

in a Fe(Ill} frameworlc}

OH ' / " / Me Me / "-if "- o mr

"\..,p/

OH cl "\..,O '\..,/'\..,/

Me Me / "-o/ "oH

OH 0 '\. / " / '\. Me Me Cr-OH / "-o/ tn{

Surface complex for-mation to bi- or multi-nuclear complexes or

surface films, blockage of surface gr~ups by

metal cations

Schematic representation of factors that promote or inhibit dissolution of hydrous oxide.

b)

OH /tj:I ' /' ' _Jlf Ffl-11

Organic Reductant

e.g., Ascorbate

Fe- / b /b H~ast

w HO 0 OH 0 t;/t fastj '\,~cf '\.. / '-Feil( Y

Hj( -a

-o/ ~o / 'r:/ Fem "-a"'-o !Electron H OH 0 ~ Transfer H OH /'\. /'\..,c 01\ / '- .,in Fem C OH

\.,i{ Fe'"d /F 'o/ V ~o '.,il "-Fen__ Ji / '-68 JI I slow J V ls low t I

0 H ' /0

H I / <; 0 H slow H ~O'\. 70) f / H m Fe'" + Fri /C~ ' / ·f;l I -Fell '. e'" + Fee,,, / '. ".0 ~ C n Fem Fe:.,, ''

/ '-oH OH / 'Q{j'

a) The structures given here are schematic short hand notations to illustrate the principal feature. Charges given are relative. b) exemplifies schematically pathways for enhanced dissolution by acids (H+ ions), ligands (example oxalate), and reductants (example ascorbate). In each case a

surface complex (proton complex, oxalato and ascorbato surface complex) is formed, which influences the bonds of the central Fe ions to 0 and OH on the surface of the crystalline lattice, in such a way that a slow detachment of a Fe(III) aquo or a ligand complex (in case of reduction an Fe(II) complex) becomes possible. In each case the original surface structure is reconstituted, so that the dissolution continues (steady-state condition). In the redox reaction with Fe(III), the ascorbate is oxidized to the ascorbate radical A-;-. The principle of proton-promoted and ligand-promoted dissolution is also valid for the dissolution (weathering) of Al-silicate minerals. The structural formulas given are schematic and simplified; they should indicate that Fe(III) in the solid phase can be bridged by 0 and OH.

25

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

(31)

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, the ligand-promoted and reductant-promoted dissolution in terms of plausible simple reaction steps is illustrated in Figure 16b.

The dissolution rate is actually the sum of all the different parallel dissolution rates of the various metal conters. An inhibitor may compete for available surface sites.

Inhibiting Ligands. Some ligands like oxalate (and other dicarboxylic acids), salicylate, F-, EDTA (Ethylenediaminetetraacetate ), NT A (Nitrilotriacetate) accelerate the dissolution, while others, Croi-, benzoate tend to inhibit the dissolution. Some like phosphate and arsenate, accelerate the dissolution at low pH and retard dissolution at pH> 4. How can these different effects of ligands be explained? Why do similar kinds of bidentates such as CrOJ" (inhibitor) and oxalate (dissolution promoter) act so differently? One hypothesis is that mononuclear complexes, especially if they are bidentate (and mononuclear), accelerate dissolution, while binuclear complexes that bridge two surface functional groups inhibit dissolution. The inhibitive effect of binuclear complexes could plausible be accounted for by inferring that much more energy is needed to remove simultaneously two center atoms from the crystalline lattice. In case of phosphate and arsenate one could postulate (and there is some supporting experimental evidence) that at low pH mononuclear and at near neutral pH value bi- or tri-nuclear surface complexes are formed. In case of CrOf and oxalate (C2ol-) one may imply that in the former case binuclear and in the latter case mononuclear complexes (both bidentate) are formed. The electron donor properties of the (oxidizing) chromate and (potentially reducing) oxalate ligands are remarkably different. A further role of the cr0i- is to maintain a high redox potential at the oxide-water interface and thus to impart passivity and to prevent reductive dissolution. Any chromate that is reduced becomes adsorbed, even at relatively low pH values, as Cr(ill) to the oxide surface; the adsorbed cr3+ or CrOH2+ surface complex is an excellent inhibitor for the dissolution.

As shown in Figure 17a, phosphate and borate inhibit the reductive dissolution of goethite by H 2S. Similarily, the non-reductive dissolution of lepidocrocite (y-FeOOH) by EDTA (Y4-) is inhibited by phosphate and arsenate (Figure l 7b). Both in the reductive dissolution (by H2S) and the ligand-promoted dissolution (by EDTA), the inhibition effects can be explained by ligand-exchange reactions of the type (the definitive assignment of the protonation of surf ace species has not yet been established).

2 =FeS- + H2P04 + 3 H+ = (=FehHP04 + 2 H2S 2=FeH2Y-+H2P04+H+ = C=Fe}iHP04 +2H3Y-

(32a) (32b)

26

0 100 200 time [min]

c)

Figure 17

300 0 10

8 !

without Cr{]II}

0.001 M Cr( ID)

0 ---~--.-~----...~----' 0 20 40

time[bours] 60 80

20 time [hours]

30

Inhibition of the dissolution of Fe(III)(hydr)oxides a) Effect of phosphate and borate on progress of reductive dissolution of goethite by H 2S. (10-2 M NaCI04)

pH= 5.0. Partial pressure of H2S = 10-3 atm. 0.032 g .e-1 goethite. b) Effect of phosphate and arsenate on retarding the EDTA 00-3 M)-promoted nonreductive dissolution rate of

lepidocrocite (0.5 g .e-1) at pH= 7.0 (0.01 M NaC104). c) Effect of 10-3 M Cr(Im on the proton-promoted dissolution (pH= 3) of a-FeOOH (0.5 gl l) 0.1 M KN03•

(From Bondietti et al., 1993)

40

X-Ray Studies on Surface Complexes of Oxoanions on Fe(IIl)(hydr)oxide Surfaces. The local coordination environment of arsenate absorbed on and coprecipitated with ferrihydrite was investigated using extended X-ray absorption fine structure (EXAFS) spectroscopy (Waychunas et al. 1993 and Manceau, 1995). The presence of a bidentate-binuclear complex was noted by both Waychunas et al. and Manceau and appears to be the dominant surface complex. The same type of complex was also observed for phosphate and arsenate sorption on goethite using infrared (IR) spec-troscopy. Recent wideangle X-ray scattering (W AXS) data lend further support to the bidentate-binuclear complex (Waychunas et al., 1996). Recent EXAFS studies confirmed the existence of binuclear surface complexes of arsenate and chromate on the goethite surface (Fendorf et al., 1997). Partially binuclear binding of phosphate by Fe oxohydroxides has also been suggested on the basis of EXAFS studies by Rose et al. (1997).

It seems justified to postulate that a passive oxide film reacts sirnilarily as a hydrous oxide with regard to acids and ligands. The dissolution rate of the passive film would thus depend on H+ and other solution variables. Figure 18 represents schematically the reactivity of the dissolution behavior of a iron metal covered by a passive oxide film.

27

Film of Metal Fe( ilJII lihydr)Oxide

I H

-HO OH '-./

Fd!I), FeOII} l I /' ~-0 OH

: I L-... Fe(Ill)+e-

1

e----....

Figure 18

Electrolyte

Ligand (orinlribibr) H+

fe(ll,ill)(aq)

02

Hp

Idealized representation of a passive Fe surface. A steady state between the dissolution of Fe (corrosion) and the passivating hydrous Fe(II,III)-oxide film is assumed. The dissolution of the oxide film is influenced by dissolution promoting acids (W) and ligands (e.g., F-, oxalate, ascorbate) or inhibitors (borate, phosphate, oxidants) in the electrolyte. The oxide film may be electron-conducting.

Figure 19 compares the pH dependent dissolution rate of various minerals with that of various passive films. The latter were obtained from the anodic current in the passive region of poteniostatic polarization curves. It is interesting to note that the dissolution rates of oxide films are two or three orders of magnitude faster than that of oxide minerals, While a similar pH dependence of the rate is observed. No systematic study on the effects of ligands and inhibitors (examining separately the effects of pH and ligands) have been made, although semiquantitative evidence supports the idea that bidentate ligands (such as oxalate) and high pH (above 11) accelerate and that phosphate and borate retard the dissolution rate of Fe oxide film. Polarization curves (cyclic voltarrunograms) on pure iron (99.99 %) in acetate buffer (pH = 4.6) illustrate that the passive current increases with oxalate but decreases with benzoic acid, phosphate and borate (Sinniger, 1992).

The oxoanions borate, phosphate, arsenate, selenite, and silicate which have been shown (Figure 17) to inhibit both non-reductive and reductive dissolution of Fe(ill)(hydr)oxides, are known to enhance the passivity of oxide films (Szklanska-Srnialowska, 1993).

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); borate has been identified to be present in the film (MacDougall and Graham, 1995). The oxoanions, in addition to their ability to inhibit dissolution, imparts on the iron(hydr)oxide surface a negative surface charge in the neutral pH range. The negative surface charge imposed by the surface complexes of oxoanions may retard the migration of er and of dissolution promoting ligands and reductants to the Fe surface (Strehblow, 1995). Oxidizing oxoanions like CroJ- and MoOf act both as dissolution inhibitors and oxidants maintaining high (passivating) redox potential and preventing reductive dissolution. Any Cro~- that is reduced becomes adsorbed as Cr3+ or CrOH2+. Adsorbed Cr(ill) is also an efficient dissolution inhibitor (Figure l 7c).

28

1(}7

v: N

·= § JQ-9 (!) -;; .... ~2 ] ~ 10-11 v. :0

10-13

Figure 19

2 4

minerts \ l

\'calcite \' \:------_

Doloirlte \

6 pH

8 10

passi~films

ID-4 I j 1(}6 - r i

105 ~e IO/, .5 M S)4) ~ lfrl

t(}7 - ~ \'\ I ; ·~Fe(0.15M, 1

106 f- \ phrnphate) -i lQ-2 , ' I I \ Co (pho41hate.

107 I \ ~fe)-l J(}3

1(}8 -

}(}9 - I\" ~i (05 M 00 2} ! 1()8 l- \ 4

. 104 l()·lO !Cr (Hp04) /

}(}9 L Ti (SO~·. Ac-) j I0-5

1(}11 - I ]()10 ~

10-12 -I

1()11

}(}13 -

0 2 4 pH

6 8

Comparison of dissolution rates of minerals (a) and of passive oxide films (in parentheses, electrolyte) (b ). Data for Figure a) are referenced in Stumm ( 1997). In Figure b) the lines are sketched from data obtained by potentiostatic anodic polarization curves (current density in the passive region) referenced in the review by Sato and Okamoto ( 1981 ). A steady state of passivity maintaining current is assumed (rate of metal dissolution equals the dissolution rate of passive film). For the conversion of passive current into dissolution rate it was assumed that one mol of metal is dissolved per two Faradays (F) (dissolution rate [mol m·2 s·11 = i [A cm"2] x 104/2 F).

Corrosion Pits. Ligand promoted dissolution of the oxide films assists in the breakdown of passivity and promotes the formation of corrosion pits.

Experiments by Strehblow ( 1995) with the use of the rotation-ring-disc electrode on the effect of fluoride (0.1 M F-, low pH) show that :fluoride-promoted Fe(ill) oxide dissolution proceeds as follows (using our notation for surface species,

(33a)

(33b)

29

fast FeF(aq)2+ + n HF ~ FeF n(aq)3-n + n W (33c)

where the detachment step (29b) is rate determining. The same surface controlled mechanism has been proposed by Zutic and Stumm (1984) 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 aluminum electrode. Although er and so~- form weaker complexes with Fe(ill) than F-, a similar mechanism has been proposed for the aggressive action of soJ- and chloride at low pH.

1. 8 Reduction of Chlorinated Solvents such as Carbon Tetrachloride and Hexachlorethane at a Corroding Iron Surface

Corroding zero-valent iron is not only a reductant of H+ and 0 2 but does also reduce organic substances such as chlorinated hydrocarbons, RCL

Fe0 + RCl + H+ --.. RH + Fe2+ + er (34)

i. e., the Fe0 is oxidized by RCI.

The electrochemical parameters of this reaction were studied with the help of a rotating disk electrode in an electrochemical cell under controlled conditions (pH= 8.4, borate buffer) (Scherer et al., 1997). Under the conditions of this study the rate of reduction of CC14 by bare Fe0 appears to be dominated by the reaction at the metal-water interface rather than by transport to the metal surface; the electron transfer at the iron surface controls primarily the dechlorination rate.

Figure 20 gives the systemic mixed potential diagram showing j-E curves for the Fe0-CC14-H20 system derived from experimentally determined kinetic parameters. The data were obtained under conditions where the iron is not covered by an oxide layer.

At the corrosion potential, EcorP the net current density is equal to zero, the partial current densities are non-zero; the corrosion current density, icorP is given by

(35)

where jFe' fac1 and fo are the partial current densities associated with the respective rate determining steps in the oxidation of Fe0 (Fe0 ---. Fe2+ + 2 e-), the reduction of RCl (RCl + H+ + 2 e- ---.. RH + Cr), and of reduction of H20 or H+ (2 H20 + 2 e- ---. H2(g) + 2 OH-, or 2 H+ + 2 e- ---. H2(g) ). The rate ofRCl reduction can be calculated from Faraday's law

(36)

where n is the number of electrons in the reaction, F is the Faraday constant and C~Cl the concentration of RCl at the surface; kct is the rate of dechlorination reaction (charge transfer) which is dependent on the electrode potential.

Table 3 gives the kinetic parameters that characterize the factors that regulate the electrochemical reduction of CC14 by Fe0.

30

Table 3 Electrochemical Kinetic Parameters for Fe0-CC14-H20 System (from Scherer et al., 1997)

Redox Couple Eiq aJ (V) j~ (A cm-2)

Fe0 --. Fe2+ -0.620 1.5 x 10-6

2 H+--. H2(g) -0.500 5.5 x 10-6

CC14 --. CHC13 +0.720 2.2 x 10-9

Px (V) (Tafel Slopes)

0.077 0.061 0.127

a) Equilibrium potentials were estimated using the Nernst equation for the conditions pH 8.4, [CC] = [Fe2+] = 10-6 M, [CC14] = [CHC13], and pH2 = 1 atm.

Figure 20 can be constructed by considering the Tafel equations for the partial current densities.

where j~ (for reaction x) is the exchange current density, Ex is the electrode potential, Eiq is the equilibrium in potential for the half reaction (see Table 3), Px is the Tafel slope; positive signs are inert for anodic current, and negative for cathodic current.

As seen in Figure 20 at Ecorr (zero net current), j 8 is negligible compared to jRCI; thus, JRCI :::::: icorr· RCl is the cathodically reduced oxidant in this corrosion reaction, a role usually played by oxygen in conventional aerobic corrosion.

30

20 :icorr I

10 Fe0-Fe2+ I I I

c.<i"" "' I

·a :E u 0

:/ corr

< ~ -0.6 -0.4 EH (V)

x ._.., -10

-20 ~

2 H+-H2(g)

-30

Figure 20 Reduction of CC14 by corroding Fe. Mixed potential diagram showing partial current densities vs. E curves based on electrochemical parameters determined experimentally on a rotating iron electrode in an electrochemical cell.The corrosion rate reflected by icorr is balanced by the reduction rate of H+ or of H20 (almost negligible) and the reduction rate of CC14• The diagram is valid for the redox reaction at a bare (oxide-free) iron surface in presence of a borate buffer (pH = 8.3). (From Scherer et al., 1997)

31

In Presence of an Oxide Film. The reduction of CC14 has also been investigated (Johnson et al., in print) in presence of an oxide film on the iron surface. Despite the presence of the oxide film, re-duction of CC14, although slower, is still possible. The chlorinated hydrocarbon is assumed to be-come adsorbed as an outer-sphere complex to the oxide, where it becomes reduced, i.e., the redox reaction with the Fe0 is mediated by the oxide. The dehalogenation is inhibited by borate which en-hances passivation (by rendering the oxide film less dissolvable) and accelerated by chloride which destabilizes the passive film. Dehalogenation is also inhibited by catechol and other ligands that may block the adsorption of CC14 to the active sites of the oxide film.

1.9 Metals other than Iron

The principal materials used in water supply distribution systems are cast iron, ductile iron and steel. These materials are distinguished based on differences in their non-ferrous portions (Table 4).

Table 4 Typical Composition of Iron Materials

Sample Elemental Composition % c Si Mn p s Ni Fe

Ductile 3.7 2.5 0.45 0.03 0.01 0.6 92.7 Cast 3.3 2.2 0.6 0.20 0.12 93.2 Steel 0.15 0.17 0.5 0.03 0.03 99.1

Plumbing materials other than Fe used to transport drinking water are Zn (galvanized pipe) Cu and Pb.

Galvanized Pipe (Zn) Zn is more anodic ("active") than Fe (Table 1). In Zn-coated iron metal, the Zn acts as an anode and dissolves there by protecting the iron cathodically (Chapter 1.1). Figure 21 illustrates some experiments with an aggressive water (Bachle et al., 1981). The change in the composition of the scale layer, early in the pipe's life, Zn-compounds were the principal components, but after the Zn-coating vanished corrosion of iron became dominant.

32

600 AZn eFe

500 NE

200 e Oh

i:: N

,,,; t>

0£) 400 ca c ·:g 0 u

::::s

'8 a c 0 ·;:;;

u 300 = .i:S ....

0

g 8 .... 100 0

"' "' Cl.) = 200 ~ u :E

c • ~ c 0 u

E--< c ,g 100

0 800 0

200 400 600 Time, days (with galvanized)

0 200 400 600

Time, days (without galvanized)

Figure 21 Rapid loss of zinc in an aggressive water. Comparison of iron content and Zn content of corrosion products on galvanized steel. (Source: Bachle et al., 1981)

100

10-2 ,......... :E ...._,, i:: 10-4 0 ·-~ .::; i:: a)

10-6 u i:: 0 u

10-B

Figure 22 pH-dependence on solubility of ZnO(s). The solubility is given by TOTZn(ll) = [Zn2+] + I, [Zn(OH)i°n]. (Adapted from Stumm and Morgan, 1996)

33

Figure 23

0

2

4

gi 6 I

8

10

2

malachite Cu2(0H)2C03(s)

4 6 8 pH

tenorite CuO

IO 12 14

Solubility of Cu(Il) as a function of pH for Cr ([H2C0*3] + [HC03] + [C02-]) = 10-2M. The solid line surrounding the shaded area gives the total solubility of Cu(U), which up to a pH value of 7 .0 is governed by the solubility of malachite [Cu2(0H)iC03(s)J. In the low pH region, azurite [Cu3(0H)2(C03Ms)] is metastab1e but may become stable at higher CT. Above pH 7, the solubility is controlled by the solubility of CuO (tenorite). The predominant species with increasing pH are Cu2+, CuC03(aq), Cu(C03)~-. and hydroxo copper(II) anions.

From Stumm and Morgan ( 1996) where the equilibrium constants used in the construction of the diagram are given

The rate of corrosion of Zn is retarded by a coating of ZnO(s) or Zn(OHh(s) formed in natural wa-ters. Thus, the solubility of ZnO(s) gives some ideas on how pH affects the corrosion of Zn (Figure 22).

Copper In aquatic oxic systems, copper is usually coated with Cu20(s) (cuprite) and CuO(s) (tenorite) or Cu(OHh(s) and possibly by solid hydroxi-carbonates such as malachite (Cu2(0H2C03(s)). If corro-sion were uniform no pipe failure should occur but under low pH conditions Cu(II) concentration in the drinking water may exceed tolerance concentrations. The solubility of Cu(II) for a carbonate bea-ring water as a function of pH is given in Figure 23; it is obvious that below pH = 7, Cu(Il) can ex-ceed concentrations larger than 1 mg/ e. The other problem is pitting that may occur as a consequence of heterogeneities on the surface of the metal.

Various causes for such inhomogeneities can be given, for example: 1) residues of mill scale that have not been removed prior to the installation; 2) localized carbon films that remain from drawing of the pipe; 3) anomalies in the oxide film; and 4) irregular hydraulic flow patterns.

34

The development of pitting in copper tubes caused by such inhomogeneities has been discussed by Edwards et al. (1990).

Lead Corrosion of lead is of concern not because of pipe failure but because it produces a toxic corrosion product. The World Health Organization (WHO) set the acceptable level of lead in drinking water (1993) to an average value of 10 µg/liter (5 x 10-8 M). Although lead pipes are in many countries no longer in use, Pb is often used in distribution systems as leaded solder1 >. Tin-lead solders are anodic to copper but cathodic to steel.

In natural waters above pH= 5 lead is covered by lead carbonate or by Pb3(C03)z(OH)z(s) (see pe vs pH diagram Figure 24a) or in absence of carbonate by PbO(s). The solubility of Pb(Il) as a function of pH is given in Figure 24b) for Cr = 10-3 M. As these figures illustrate, in low alkalinity waters and pH < 8, toxic concentrations of Pb2+ may result.

a) b) 0

-1 -2

1.2 +20 -3 -4

-5 g -6

+15

~ -7 .I:l !::. -8 bQ .9

-9

0.6 +10

€ w ::i:: c..

I.LI +5 -10

0 Pb(OH)3 -11

-12

-0.3 -5 -13

-14 5 10 3 4 5 6 7 8 9 10 11 12 13 14

pH pH

Figure 24 a) pe (E8 ) vs pH diagram for the Pb, H20, C02 system. TOTPb = 104 M, TOTC03 {[H2C03] + [Hco;J +

[COi-D = 10-2 M (25 °C) passivity may exist in the shaded areas. {From Stumm and Morgan, 1996)

b) Log [Pb2+] vs. pH for TOTC03 = 10-3 M and 25 °C as controlled by three different solids: PbO, PbC03 and Pb3{C03)z(OH)2. Unit activity coefficient in the aqueous phase and unit activities for the solids are assumed. The solubility given for PbO(s) corresponds to the theoretical solubility of Pb2+ in absence of aqueous carbonate. {Adapted from Pankow, 1991)

1) Solders are typically used mainly to join copper pipes. {Nonleaded solder compositions are available for potable water applications.)

35

Cement-Mortar lined Pipes (Cast Iron or Steel) Linings are applied either in manufactoring plants or in situ. The latter method is used to rehabilitate water mains.

Cement-mortar lining has been successful of preventing corrosion in water mains that gave problems for unlined systems (corrosion and red water). A further advantage is that the linings ensure high flow hydraulic carrying capacity. Usually, degradation of cement materials is slow; high so~- con-centrations (> 3 mM) however, causes rapid degradation, because so~- reacts with A120 3 compound in the cement to form a highly hydrated calcium aluminum sulfate mineral.

36

Reading Suggestions

Landolt, D. Corrosion et Chimie de Suiface de Metaux, Presses Polytechnique et Universitaire Romandes, Lausanne, 1993

Bockris, J. O'M. and Khan S. U. M. Swface Electrochemistry, Chapter 8, Plenum Press, New York, 1993

American Water Works Association: Internal Corrosion of Water Distribution Systems; Cooperative Research Report 2nd Ed., Denver, CO, 1996

Stumm, W. Chemistry of the Solid-Water Interface, Wiley-Interscience, New York, 1992

37

References

Bachle, A., Deschner, E., Weiss, H. and Wagner, I., Corrosion of galvanized and unalloyed steel tubes in drinking water of different hardness and increased neutral salt content, Werkstoffe und Korrosion 32, 435-442 (1981)

Biber, M. V., Dos Santos-Afonso, M. and Stumm, W., The coordination chemistry of weathering: IV. Inhibition of the dissolution of oxide minerals, Geochim. Cosmochim. Acta 58(9), 1999-2010 (1995)

Blesa, M., Morando, B. J. and Regazzoni, A. E., Dissolution of Metal Oxides, CRC Press, Boca Raton, FL ( 1994)

Bockris, J. 0. M. and Reddy, A. K. N., Modem Electro Chemistry, Plenum Press, New York (1970)

Bondietti, G., Sinniger, J. and Stumm, W., The Reactivity of Fe(ill)(hydr)oxides; Effects of Ligands in Inhibiting the Dissolution, Colloids and Surf. A 79, 157-167 (1993)

Edwards, M., Ferguson, J. F. and Reiber, S. H., The pitting corrosion of copper, J. Amer. Water Works Ass. 7, 74-90 (1994)

Fendorf, S., Eick, M. J., Grossl, P. and Sparks, D. L., Arsenate and Chromate Retention Mechanisms on Goethite. 1. Surface Structure, Environ. Sci. Technol. 31(2), 315-320 (1997)

Franck, U. F. and Weil, K., Zur Korrosion des passiven Eisens in Schwefelsaure, Z. Elektroch. 56, 814 (1952)

Furrer, G. and Stumm, W., The Coordination Chemistry of Weathering: I. Dissolution Kinetics of 8-Al203 and BeO, Geochim. Cosmochim. Acta 50, 1847-1860 (1986)

Grossl, P. R., Eick, M., Sparks, D. L., Goldberg, S. and Ainsworth, C. C., Arsenate and Chromate Retention Mechanisms on Goethite. 2. Kinetic Evaluation using a Pressure-Jump Relaxation Technique, Environ. Sci. Technol. 31(2), 321-326 (1997)

Hering, J. and Stumm, W., Oxidative and Reductive Dissolution of Minerals, in Reviews in Mineralogy 23, M. F. Rochella Jr. and A. F. White, Eds., Mineralogical Society of America, Washington, D. C., 427-465 (1990)

Heusler, K. E., Iron, in Encyclopedia of Electrochemistry of the Elements, Vol. 9, Part A, A. J. Bard, Ed., M. Dekker, New York, 229-381 (1982)

Johnson, T. L., Fish, W., Gorby, Y. A. and Tratnyek, P. G., Degradation of Carbon Tetrachloride by Iron Metal: Complexation Effects on the Oxide Surface, J. Contaminant Hydrology (in print)

Landolt, D., Corrosion et Chimie de Surface de Metaux, Presses Polytechnique et Universitaire Romandes, Lausanne (1993)

MacDougall, B. and J., G. M., Growth and Stability of Passive Films, in Corrosion Mechanisms in Theory and Practice, P. Marcus and J. Oudar, Eds., M. Dekker, New York (1995)

Manceau, A., The mechanism of anion adsorption on iron oxides; evidence for the binding of arsenate tetrahedra on free Fe(O,OH)6 edges, Geochim. Cosmochim. Acta 59, 3647-3653 (1995)

Morsi, M.A., Elevadi, Y. A., Lorbeer, P. and Lorenz, W. J., The Inhibition of the Iron Corrosion in Aqueous Solutions Containing Oxygen, Werkstoffe und Korrosion 31, 108-114 (1980)

Pankow, J. F., Aquatic Chemistry Concepts, Lewis, Chelsea, MI (1991) Pou, T. E., Murphy, P. J., Young, V. and Bockris, J. 0. M., Passive Films on Iron: The

Mechanism of Breakdown in Chloride Containing Solutions, J. Electrochem. Soc. 131(6), 1243 (1984)

38

Rose, J. M., Flank, A. M., Mason, A., Bottero, J. Y. and Elmerich, P., Nucleation and Growth Mechanisms of Fe Oxyhydroxides in the Presence of PO 4 ions, Langmuir 13, 1827-1834 ( 1997)

Sato, N. and Okamoto, G., Electrochemical Passivation of Metals, in Comprehensive Treatise of Electrochemistry, J.M. Bockris et. al., Eds., Plenum Press, New York (1981)

Scherer, M. M., Westall, J. C., Ziomek-Moroz, M. and Tratnyek, P. G., Kinetics of Carbon Tetrachloride Reduction at an Oxide-Free Iron Electrode, Environ. Sci. Technol. 31(8), 2385-2391 (1997)

Schindler, P. W. and Sposito, G., Surface Complexation at (Hydr)oxide Surfaces, in Interactions at the Soil Colloid-Soil Solution Intelface, G. H. Bolt et al., Eds., Kluwer, Dordrecht (1991)

Sinniger, J., Die Inhibition der Auflosung von Eisen(III)(hydr)oxiden und ihre Beziehung von Passivitiit von Eisen, Ph.D. Thesis, ETH Ziirich (1992)

Stratmann, M. and Muller, J., The mechanism of the oxygen reduction on rust covered metal substrates, Corrosion Sci. 36, 327-359 (1994)

Strehblow, H.-H., Mechanisms of Pitting Corrosion, in Corrosion Mechanisms in Theory and Practice, P. Marcus and J. Oudar, Eds., M. Dekker, New York (1995)

Stumm, W., Evaluation of Corrosion in Water by Polarization Data, Industrial and Engr. Chem. 61, 1487-1991 (1959)

Stumm, W., Investigation on the Corrosive Behavior of Waters, J. San. Engr. Div. ASCE 86, 27 (1960)

Stumm, W., Ed., Aquatic Sulface Chemistry, Wiley-Interscience, New York (1987)

Stumm, W., The Chemistry of the Solid-Water Inteiface, Wiley-Interscience, New York (1992)

Stumm, W. and Morgan, J. J., Aquatic Chemistry 3rd ed., Wiley Interscience, New York (1996)

Stumm, W., Reactivity at the Mineral-Water Interface; Dissolution and Inhibition, Colloids and Sulfaces A 120, 143-166 (1997)

Szklanska-Smialowska, S., Inhibition of Pitting Corrosion, in Reviews on Corrosion Inhibitor Science and Technology, A. Raman, Ed., NACE, National Asociation of Corrosion Engineers, Hustontrexas (1993)

Venema, P., Hiemstra, T. and van Riemsdijk, W. H., Interaction of Cadmium with Phosphate on Goethite, J. Colloid lnteif. Sci. 192(1), 94-103 (1997)

Waychunas, G. A., Rea, B. A., Fuller, C. C. and Davis, J. A., Surface chemistry of ferrihydrite: Part 1. EXAFS studies of the geometiy of coprecipitated and adsorbed arsenate, Geochim. Cosmochim. Acta 57, 2251-2569 (1993)

Waychunas, G. A., Fuller, C. C., Rea, B. A. and Davis, J. A., Wide angle X-ray scattering (W AXS) study of "two-line" ferrihydrite structure: Effect of arsenate sorption and counterion variation and comparison with EXAFS results, Geochim. Cosmochim. Acta 60(10), 1765-1781 (1996)

Westall, J.C., Adsorption Mechanisms in Aquatic Surface Chemistry, in Aquatic Suiface Chemistry, W. Stumm, Ed., Wiley-Interscience, New York (1987)

Zutic, V. and Stumm, W., Effect of Organic Acids and Fluoride on the Dissolution Kinetics of Hydrous Alumina. A Model Study using the Rotating Disc Electrode, Geochim. Cosmochim. Acta 48, 1493-1503 (1984)

39