An electrochemical study of phase-transitions in rust layers

Corrosion Science. Vol. 23, No. 9, pp, 969-985, 1983 Printed in Great Britain.

0010-938X/83 $3.00+0.00 ~) 1983 Pergamon Press Ltd.

A N ELECTROCHEMICAL STUDY OF PHASE-TRANSITIONS IN RUST LAYERS M.

STRATMANN,K,

BOHNENKAMPand H.-J. ENGELL

Max-Planck-Institut fiir Eiscnforschung GmbH, Dfisscldorf, FRG

Aimmct--Isolated rust layers have been investigated by electrochemical methods to find out whether their reduction and re-oxidation can affect the atmospheric corrosion of iron. At potentials below 0 mV, first a thin Fet+-containing surface layer is formed on top of the y-FeOOH. This reduced surface layer can dissolve into the cell electrolyte at acid pH, or at lower potentials the FcS+-ions can react with y-FeOOH to FcsO,. The formation of magnetite could be followed by tn-situ magnetic measurenmats. The reduced surface layer can easily be oxidized back to y-FeOOH, magnetite can partly be oxidized to y-FeaOs.

INTRODUCTION Tin3 ATMOSPm~aC corrosion of iron is an electrochemical process which can start only in the presence of an electrolyte, as a result of sufficient wetting of the metal surface, x The primary reaction products formed are hydrated Fe s+ ions which are further oxidized by oxygen to Fe s+ and then precipitated as oxides, hydroxides and oxyhydroxides (FeOOH). The thus formed rust layer covers the metal surface to a great extent and has a large influence on the further development of the corrosion process. Long time field tests have proved that weathering steels with small contents of Cu and Cr show far better corrosion resistance than unalloyed samples3 "1° Laboratory investigations n - l ' have clearly demonstrated that this difference only occurs in the case of repeated wetting and drying of the sample. Corrosion tests in an atmosphere of constant humidity (e.g. 100 %) show no differences in the corrosion behaviour between unalloyed and weathering steels. 15 U. R. Evans ~e-18 developed an electrochemical model to explain the observed influence of changing wetness on the atmospheric corrosion of iron. H e postulated that in periods of high water content within the pore structure of the rust the anodic dissolution of iron is balanced by the cathodic reduction of the Fe(III)-oxides in the rust layer: Fe ~ Fe ~+ + 2e

(i)

Fe s+ + 8 F e O O H + 2e ~ 3 FesO4 + 4 H20.

(2)

Manuscript received 11 February 1983. 969

970

M. SIRArMANN,K. BOHNENKAMP and H.-J. ENGELL

Later, after partial drying of the pore structure the magnetite is re-oxidized by the oxygen which has now free access through the pores by gas diffusion:

3 9 3 FesO4 + ~ O2 -~ ~ HgO -* 9 FeOOH.

(3)

After wetting the cycle of reduction of FeOOH and oxidation of FesO4 can start again by reaction (2). A number of authors ts'~s have demonstrated that atmospherically formed rust layers can easily be reduced. They always observed the formation of Fe304, usually from y-FeOOH, whereas during the reduction of thin dense electrodeposited 7-FeOOH layers no magnetite is formed and the total reduction current produces free Fe s+ ions: e Obviously the reaction scheme of rust reduction is not well understood so far. No experiments have been published yet on the re-oxidation of reduced rust layers. It was the goal of the present study to examine the mechanism of rust reduction and reoxidation. To keep the system as simple as possible isolated corrosion layers were used which contained neither metallic iron nor remarkable contents of Cu and Cr. EXPERIMENTAL METHOD Rust layers isolated from metallic iron have been prepared by oxidizing iron layers of about 7 Ima thickne~ which were electroplated on to a gold electrode as an electronic conductive support (A = 7 cms) from a 0.4 M (NH,)s Fe(SO~)2 solution at pH -- 6-7 and 95°C. The electroplated iron contains approximately 0.01 ~e Cu and 0.005 ~, Cr. The rate of rust formation in oxygen of 99~, r.h. with addition of 1.5 g SO~/mI Fe was followed volumetrically, aDThe rust layers were further characterized by X-ray, magnetic and chemical methods and by a BET measurement to determine their phase composition, their amount of ferromagnetic components such as a-Fe and FesO4, the amount of Fe~÷and Fe *+and the specific surface of the layers. An oxygen-free 0.2 M Na~SO4 solution at T = 20°C was used as an electrolyte. Some measure. ments were conducted in a 0.2 M NaCIO~ solution. They are not discussed in this paper as the results do not differ considerably from those of the sulphate electrolyte. The pH was kept constant (4- 0.1 pH) by regulated addition of 0.01 N NaOH and 0.01 N H~SO4. The Fe s+ concentration of the solution was determined by AAS and chemical analysis with o-phenanthroline. The electrochemical arrangement comits of a fast-rise potentiostat (Wenking, model POS 73) and sensitive integrator (Metrohm). The Pt counter electrodes were separated from the compartment of the cell containing the sample by a diaphragm. The reference electrode was a Hg/HgSO,-electrode. All potentials in this study are related to the standard hydrogen electrode (SHE). It has been proved that a background current due to electrochemical processes at the gold electrode was lower than 0.4 IrA cm-I within this electrolyte and the potential region of the measurements. Thus, it does not influence the accuracy of the results. Magnetite being a crucial product within the sequence of the reactions, an arrangement was developed which allowed in situ the formation and oxidation of this oxide to be followed by measuring the force, acting upon the sample within an inhomogeneous magnetic field. The sample is suspended at a sensitive electronic balance (Sartorius, model 1265), the electrical connection of the sample to the potentiostat is made by the suspension. The field strength could be varied between 200 and I000 Oe. Supplementary e x situ measurements were performed at 8000 Oe. At 8000 Oe the force is proportional to the amount of FeaO4 and a calibration was possible by means of a compact as well as of a flue

grained FesO,, prepared by partialoxidation of a Fe(OH)z-suspension by KNOB. A correlationof the in situ data at 200 Oe and the e x situ measurements at 8000 Oe is possible, thus obtaining the amount of FeaO, during the polarization. The detection limit of the magnetic measurement is about 0.I-0.3 rag Fe~04. In the following sections the measured values of electrical charge and magnetic force always are divided by the iron content of the rust layer to allow a comparison between the results of different samples. Usually mr, is around 35 ms.

Pha~-transitions in rust layers

971

EXPERIMENTAL RESULTS AND DISCUSSION

Characterization of the rust layers Figure 1 shows the oxygen uptake of the samples during the atmospheric corrosion related to the amount of eleetroplated metallic iron. After about 150 h the iron is completely oxidized and the ratio of oxygen to iron (Ano,/nv,) reaches 0.75. That means that all the iron is oxidized to the oxidation state 3 q-. An X-ray analysis of the rust layer showed nearly equal amounts of y-FeOOH and ,,-FeOOH. Fe~O4 could be detected only in traces; lines of~-Fe were absent. Magnetic measurements confirmed that the rust layer contained nearly no ferromagnetic products such as FesO4, y-Fe2Oa, 5-FeOOI-I and =-Fe. SEM investigations indicated a very small crystal size and high porosity of the formed rust layers. The specific surface is quite large, a BET-measurement yielded a value of about 51 m a g-1 rust. So the surface of the layers examined in this study is nearly 3 m t (geometric surface: 2 cm z) and the mean crystal size of the FeOOI-I crystal should be about 40 nm (length of the edge of an assumed cube).

Reduction of the rust layer The possible reduction reactions of the FeOOH species such as reductive dissolution, u formation of FesO, by reduction, formation of FesO4 by chemical reaction with FeS+~, u and the Evans reaction (2) x8 depend in their kinetics in different manner on the parameters potential, Fe s+ concentration and pH value. So the reduction of the corrosion layer was analyzed as a function of these parameters. In each case one was changed and the other two were kept constant in the bulk of the electrolyte. Reduction as a function of potential (pH -- 6.0; [Fe s÷] = 10-6 mol 1-1). Figure 2 shows the results of a potentiostatic reduction of the rust layer at potentials lowered step by step between 0 and -- 400 mV(SHE). At -- 200 and -- 300 mV(SHE) the A noz

ne. 0.6

O.t

0.2

][

--35 m0 cT ml

IV

msoz= 0.03 mg/mgFe

I

[oo;oo.

0 100

'200 t/h

Fro. 1. Extent of oxidation of the electroplated iron layer during atmospheric corrosion.

972

M. STRATMANN,K. BOHI~mNKA~Pand H.-J. ENGELL

U/mVs~ -J NA

-200

-300

k~,0.5

j,~lO

- 400

Fmagn

j.~ZO mA-cm "2

cmz

mgre

20 ] x

~pH

[Fe;~'1: 10-6tooll-~ 60 =

,

10

~x

50.

150

100

t/h FIG. 2.

Current density ( ) and magnetization (x) of a rust layer during a potentiostatlc reduction at potentials between0 and - 400 raV(B~).

current first increases but then decreases sharply with time and almost no change in the force acting upon the sample due to the magnetic field could be observed. At -- 400 mV a considerable current flows for about 100 h and the change in magnetic force indicates in accordance with X-ray data that magnetite is formed. However, even at -- 200 and -- 300 mV the electric charge is quite large according to the high current densities at the very beginning of the reduction (Fig. 3). If the newly prepared sample is polarized in one step from the rest potential [~ + 200 mV(SHE)] to -- 400 mV even at this potential the formation of FesO4 starts only after about 1/3 of the charge for total reduction has passed (Fig. 4). As this plateau is observed for a magnetic field strength reaching from 200 to 8000 Oe it is unlikely that in this first stage of the reduction small superparamagnetic FesO4 crystals are formed because at 8000 Oe crystals of (2 nm) s should clearly be detected. By chemical analysis it could be further shown that only at the end of the reduction a small amount of Fe 2+ diffuses into the bulk electrolyte of the cell (Fig. 4). According to X-ray data only y-FeOOH is reduced during the electrochemical polarization. The amount of~-FeOOH in the rust layer is not influenced by the polarization (Fig. 4). The magnetic data of Fig. 4 can be analyzed quantitatively by e x situ calibrations using Fe304 powder compounds (see Experimental Methods) at 8000 Oe to determine the amount of FesO4 as a function of the cathodic charge. By comparing the charge and the amount of FesO4 formed it is possible to evaluate the amount of the Fez+ species which is being formed at the beginning of the reduction process before the formation of FesO4 starts (Fig. 5).

Phase-transitions in rust layers O/F

I o

I

I

I

973

I

Fmo~n 111gfe

I'lFe

3

[Fe 2'] = 10.5 tool I "1

0.10 O

pH -- 5.0

005-

-1 x

Frnogn'~ 0.00

I

- 4O0

FIG. 3.

%~'-0 . . . . ]

•

-200

0

I

0 .U

mv~

Cathodic charge ( x ) and magnetization (©) at the end of the potentiostatic reduction.

This Fe ~+ species seems to be an intermediate which is partly consumed during the magnetite formation. In this study the chemical nature of the Fe 2+ intermediate could not be determined by a X-ray analysis, It can only be stated that the intermediate is not a well-ordered phase and that the Fe 2+ ions are highly immobile as they cannot be detected in the bulk electrolyte during the major part of the reduction. Reduction as a function o f the Fe ~+ concentration (pH = 6.0; U = -- 400 mV(SHE). The Fe 2+ concentration of the bulk electrolyte has considerable influence upon the formation of magnetite (Fig. 6). In a 2 x 10 -s M Fe 2+ solution the magnetite formation starts at the very beginning of the polarization; the amount of the Fe 2+ intermediate formed is less than it was in the case of the Fe~+-free solution. The rate of FeaO4 formation is considerably higher in the FeS+-containing solution and approx. 45 % more magnetite is formed during the reduction whereas the amount of charge is approx. 15 % less. Therefore it can be concluded that during the polarization Fe 2+ from the solution is built into the lattice of Fe~O4. This could be confirmed by a chemical analysis of the Fe 2+ content o f the sample. Finally, the potential at which the Fe~O4 formation starts is a function of the Fe ~÷ concentration. In Fe2+-free solution the formation began at -- 400 mV(SHE); in 2 × I0 -a M Fe 2+ solution the formation starts already at -- 300 mV(SH.E). Reduction as a function o f p H (U = -- 400 mV(SHE); [Fe ~+] ~ 10 -6 mol 1-1). Figure 7 shows the results of a reduction at -- 400 mV and p H values from 5.0 to 9.0. The amount of FeaO4 as a function of the cathodic charge shows nearly the same behaviour in the first stage of the reduction at p H values from 5.0 to 9.0 but the current density at p H = 5.0 is larger than those at pH = 6.0 and 9.0. Less magnetite is formed at p H = 5.0 and a large amount of Fe e+ ions diffuses into the bulk solution. This effect is strongly enhanced during a reduction at pI-[ = 4.0 (not shown in Fig. 7). In the latter case the FesO4 formed first is further reduced during polarization at

974

M. STRATMANN,K. BOHNENKAMPand H.-J. ENGELL I

I

I

n(FeZQ~ ) U =- z,O0 mrs, [ [ F e ; ' ] = l O "6 tool I "~ pH • 50

-O/F

Inlenslty of x-rOy braes

]]1 /

a-FeOOH ~,- FeOOH

f@]O4

l[

m

w

--

Ill

m

--

s

~ e ~ T ~210

/

/I,9011 j,

~=,.o,~

/ v ....

: ::'.°o7

11""° -03 02 0.1. 0 ~lFe

FIG. 4.

Magnetization ( x ) and amount of dissolved Fe'+ (A) as a function of the cathodic charge during reduction at - 400 mV(sn~).

400 mV to free Fe 2+ ions. At p H = 9.0 no Fe ~+ diffuses out of the rust layer (Fig. 7), the amount of FesO4 is reduced too, and X - r a y analysis indicates the formation of =-FeOOH from 7-FeOOH. At p H = 13.0 no reduction of the samples is possible at - - 400 inV. -

-

Discussion o f the reduction reactions The experiments proved that in a first reduction stage a Fe 2+ intermediate is formed and that the magnetite formation is a succeeding reaction. F r o m a thermodynamic point of view the formation of Fe 2+ is probable too as Fe *+ is the stable species at - - 400 mV and p H = 6.0 (Fig. 8). This is confirmed by the experiments of Cohen et al. sa It is very likely that the intermediate is a reduced surface layer on the y - F e O O H and perhaps also on the ~-FeOOH crystal. The Fe s+ species may be adsorbed on the large internal surface of the rust or the Fe s+ within a surface region of the F e O O H which is a few monolayers thick is reduced to Fe =+ accompanied by protonation of O s+ to O H - . This simultaneous exchange is often observed in the iron oxide/water systemA ~.3=

T-FeOOH -t- H + + e- ~ {Fe. O H . OH}. A reduced surface layer

(4)

Phase-transitions in rust layers i

._q_n

f

975

!

nr, 1

IU -- 400mVsHt,pH =60 ] [Fe ~'] : 10 "6 tool I "1

]

0.&

0.3-

02-

Ih

\ mm~

Fe~O~

• ~

o.I-

/v

. ~ . . ~ 2,

v- j

"~

Fe~n,~,dim 0

0

I00

150

200 t/h

FIG. 5. Kinetic of the formationand degrationof y-FeOOH (m), Fe*+~t~.oa. (O), and Fete4 (•) during reduction of the rust layer at - 400 mVsae. According to the experiments the surface layer can react in two ways: (1) The Fe 2+ ions may cross the solid/electrolyte interface: {Fe. O H . OH} + 2H + ~ Fe 2+ (aqu.) + 2H20.

(5)

Reactions (4) and (5) are identical to the reductive dissolution of y-FeOOH. Reaction (5) is strongly enhanced by a more acid pH and can clearly be detected during a reduction at pH = 4.0 and 5.0 (see Fig. 7) whereas this reaction seems to be not very important during reduction at pH values of 6.0 and higher. One reason for this is a pH change in the pore system of the rust layer as a consequence of the formation of the reduced surface layer by reaction (4). In this study the amount of reduction charge was measured as well as the amount of H + needed to keep the pH in the bulk electrolyte of the cell constant. During the first I0 h of the reduction only 50 ~o of the amount of H + which is consumed during reduction was added to the solution assuming that reaction (4) is the reaction of the first reduction stage. Therefore, rather alkaline pH values should result in the pore system of the rust layer and the pH of the beginning is less important for the first hours of the reduction (Fig. 7). The pH change has two consequences. First, the velocity of the reactions (4) and (5) and consequently the current density drops sharply (Fig. 2) and so the reductive dissolution is only of minor importance. Secondly, magnetite becomes the thermodynamic stable phase (Fig. 8).

976

M. STRATMANN,K. BOHNENKAMPand H.-J. ENOEI.L

mg~e5 l u:-~oo ~ .

o°

~LPH ~ ~ 0

0 o

~

o

/

2

j

;~ I x

? 2 0

510

,

I00

150 tlh

Fzo. 6. Kinetic of the Fe~O, formation during reduction of the rust layer at - 400 mV(e~) with Fe=+-concentrations of 10-~ mol 1-1 (x) and 2 x I0-s tool 1-1 (O). (2) A s a consequence of the p H change F ~ O 4 formation starts consuming y - F e O O H and partly the Fe =÷ intermediate:

Fe s+ + 27 -- F e O O H -+ FcaO 4 -[= 2H +

(6a)

{Fe. O H . OH} + 2 y - F e O O H -* Fe304 + 2H=O.

(6b)

or

This reaction can be interpreted from a crystallographic point of view in terms of a close packed framework of oxygen ions the interstices of which are occupied by relatively mobile iron ions and protons.3X, ss 7 - F e O O H has a layered structure. Within each layer O is arranged in a cubic close packing, the Fe 8+ ions occupying the octahedral holes. The layers are connected by hydrogen bridges. ~ FesO4 also is built up of a ccp O-sublattice, the Fe =+ and Fe s+ ions occupying tetrahedral and octahedral holes .5 and therefore a conversion of the 7 - F e O O H lattice to that of magnetite seems to be possible. This is confirmed by the very easy dehydration of 7 - F e O O H at about 180°C to 7-Fe=Os ss which has a lattice very similar to Fe304. s5 During magnetite formation Fe s+ ions have to move into the lattice of 7 - F e O O H and to occupy octahedral and tetrahedral holes. H + has to move out of the lattice or it is transferred from one O H + group to the neighbouring group and a layer of hydrated FeaO4 results. The high mobility of ions in the lattice of y - F e O O H is a well proved fact. s6

Phase-transitions in rust layers

977 Z.

,~ Fmog n

I

I

m~Fe

U

(Feoqu)

I

- O/F

:-400mV~E H

2/ ~ "

[FeZ'] = 104m°lld /

3

-Off -05

1,1o: 2

2,2o: 3,3o

=a.o

PH=9opH

//

Or,

13/

?

03

//

0.2

1-

///

,.°

01

2.. "i

O,

o.'os

o.ls

o.io

-0 D/F iqFe

FIO. 7. Magnetization ( ) and dissolved amount of Fe '+ (. . . . ) as a function of the cathodic charge during reduction at - 400 mV(ln~). (1, la) pH ~- 5.0; (2, 2a) pH •ffi 6.0; (3, 3a) pH ffi 9.0.

I

U~E mY

"

"

,500

~--~@

¥- FeOOH

"",.' \

\

0

a-

4

I

I

I

I

eZ'] = 10 "6mol I 1

--[

=

10-3

-

= 1.0

"

(~) a - FeOOH

"'.,,\\

Fe

a-Fe

I

I

I

6

8

10

z,

I

I

t

6

8

10

pH FIG. 8. Potential-pH diagram for the system Fe-Fe2+-FesO4-FeOOH. (a) FeOOH -= y-FeOOH, (b) FeOOH I= a-FeOOH. Data taken from Lunze" and Suzuki et al."

978

M. STZO,TMANN,K. BOHNENKAMPand H.-J. ENGELL

On the other hand, =-FeOOH has a hcp oxygen sublattice which during reduction to FeaO4 would have to be transformed into the ccp structure of this phase. The relative stability of the oxygen sublattice ensures that 0¢-FeOOH does not take part in den reaction process. The proposed reaction scheme explains the influence of the parameters Fe z+ concentration and pH: a lower pH value accelerates the formation of free Fe z+ [reaction (5), Fig. 7], whereas a high Fe z+ concentration and high pH value promote the build up of FeaO4 [reaction (6a), Fig. 6]. To test the proposed scheme further an attempt was made to transfer the Fe z+ ions formed in the first reduction stage into the electrolyte and to reduce their activity by complexation with o-phenanthroline: a7 Fe z+ + 3Phil + --~ [FePha]Z+ + 3H +.

(7)

Since during the complexation H + ions are liberated this reaction also retards the pH change in the pore system of the rust layer due to reactions (4) and (5). The decay of the current density is therefore slower and Fe~O4 formation is suppressed as Fig. 9 shows. The free Fe 2+ complex produced enters into the cell electrolyte and can be detected by chemical analysis. FeaO, is not formed and the total charge is larger by a factor of 3 than in similar experiments without complexing reagents, see Fig. 4, because now all the y-FeOOH is reduced to Fe =+, whereas during Fe3Oeformation only 1/3 of the iron-ions in y-FeOOH are reducible.

Re-oxidation of reduced rust layers Re-oxidation of the Fe =+ intermediate. If the reduction of y-FeOOH is stopped before the formation of FeaO~ starts, a reversible re-oxidation of the reduced surface

-Q/F

I

• I

nIF4 ,I

I

rife

0.5

rife

O0oVs,E -

I pH= 60

0.4

Lo.4

10.5% o- Phen. 0.3

v

0.:3

nfe z

0.2

0.2 gFe IO Fmagn

0.1 Fmogn

0

0.5

.1

i

I

I

20

40

(50

0

t/h

FIG. 9. Cathodic charge ( ), dissolved amount of Fes+ (~7) and magnetization (- - - -) during reduction at - 400 mV(s-E) in the presence of 0.5 % o-phenanthroline.

Phase-transitions in rust layers

979

layer is possible. Figure 10 shows the results of cyclic potentiokinetic reduction and re-oxidation between -- 300 and -k 600 mV(SHE). The anodic and the cathodic charge are both about 1 C/cycle and do not change very much during a great number of cycles. This is a strong indication that during these processes only surface layers are formed and reversibly oxidized. Diffusion of Fe ~+ into the cell electrolyte as well as FeaO4 formation cannot be observed during these experiments. Re-oxidation after formation of FesO4. Samples were partly reduced at potentials between --200 and --400 mV(SHE) for times between 1 and 20 h and then oxidized again by potentiodynamic polarization between the reduction potential and q- 800 inV. The rate of polarization during oxidation always was 200 mV/min.) This process was repeated several times. Figure I1 shows the observed current density-potential curves. Curves 1, 2 and 3-6 correspond to re-oxidation after reduction at -- 200, -- 300 and -- 400 mV, respectively. In the experiments 1-3, no FeaO4 was formed during reduction, for the potential was not low enough or the polarization was too short in time and the samples contained only the Fe 2+ intermediate in increasing amount. The observed oxidation peak corresponds to the oxidation of the intermediate as described before. In the experiments 4-6 an increasing amount of magnetite is formed in the samples during reduction and a part of the Fe 2+ intermediate is consumed by FesO4 formation (Fig. 5). So the first oxidation peak drops down and a second oxidation peak appears at a potential of about + 200 inV. This peak obviously corresponds to the oxidation of FesO4. During the oxidation of FeaO4 Fe z+ ions diffuse into the cell electrolyte. Magnetic measurements revealed that FeaO4 is not oxidized to the paramagnetic 7-FeOOH as proposed by Evans. In Fig. 12 the magnetization

" I ///-.

-

200

o.o

0

.200

.400

.600

U/mVsHE FIG. 10.

Potentiokineticreductionand oxidationcurve of a rust layer for the 1st and 80th cycle.

980

M. STRATMANN,K. BOHNENKAMPand H.-J. ENGELL I

J mA. cmZ

I

I

[Fe z. ] = I0 6 tool II

I

dU : 200 mY/rain dt

pH : 6.0

Ii

20-

.Fe z. 6

: I

1.0-

I

- 400

I

0

I

I

• 400

• 800

U/mVsHE

FIG. 11. Potentiokinetic oxidation curves of rust layers previously reduced at: (1) - 200 mV, for 1 h; (2) - 300 mV, for 1 h; (3) - 400 mV, for 1 h; (4) - 400 mV, for 4 h; (5) - 400 mV, for 9.5 la; (6) - 400 mV, for 15 h.

of the sample at the end of reduction and oxidation, respectively, are shown. The magnetic force which increases during reduction at -- 400 mV due to FeaO4 formation is only slightly reduced during oxidation up to + 800 mV. So during oxidation of FeaO4 the ferromagnetic properties of the crystal are maintained. The extent of reoxidation of FesO4 can be estimated from the difference of anodic and cathodic charge flown in each cycle (Fig. 12). If the reduction is stopped before FeaO4 formation starts the difference between these two values (AQ) is nearly zero (Fig. 10). If FesO4 is formed during reduction at -- 400 mV the reduction charge is always larger than the re-oxidation charge (Fig. 12). Obviously, a complete oxidation of FeaO4 is not possible.

Discussion of the oxidation reactions As discussed earlier the reduced surface layer can be reversibly reoxidized: {Fe. O H . OH} --, 7-FeOOH + e + H +.

(4a)

During potentiodynamic polarization up to + 800 mV FeaO4 can only partly be re-oxidized. During the oxidation Fe 2+ ions leave the crystal and the ferromagnetic properties are maintained. These facts suggest that the oxidation product is 7-Fe~Oa. FeaO4 and 7-Fe208 exhibit very similar crystallographic features. The 02- sublattice of both crystals is nearly indentical (ccp) and the Fe ~+ and Fe s+ ions occupy only slightly different octahedral and tetrahedral holes. 85 The formation of mixed crystals is possible. ~ The saturation magnetization of 7-Fe~Oa is only 20% less than that of FelOn. ~

Phase-transitions in rust layers I

IOIIF riFe

981

I

I

•

- 0 (Potentiostotic reduction ot -400 mVs,E)

[]

• I] (Potentiodynomic oxidotion with 200mV/min

up to ,,BOOmVsE} H

Fmogn.

mgFe

00,1 lii

t

004

2

0.020

1

0 0

5 10 number of reduction - oxidotion - cycles

15

Fit. 12. Cathodic and anodic charge exchanged between rust sample and electrolyte. Reduction at - 400 mVt~m~)for times between 1 h (lst cycle) and 20 h (16th cycle); potentiodynamie anodic oxidation up to + 800 mV(am~), sweep rate 200 mV rain-1. F~r,. = magnetic force at the end of reduction and oxidation, respectively. If the 0 2- sublattice is maintained during oxidation 2/3 of the Fe 2+ ions of FeaO4 are oxidized giving Fe s+, and 1/3 leave the crystal during oxidation to maintain electroneutrality:

3 Fe~O4 ~ 4 7-Fe203 + Fe~+ + 2e.

(8)

The extent of oxidation of FeaO4, z, may be defined as: z =

n~ (Fe30D -

n= (Fe30D

nb (FesO4)

(9)

where: nb (Fe304) = amount of FeaO4 before oxidation and n a (FesO4) = amount of FesO4 after oxidation.

As three electrons are needed to form 3 FesO4 from 9 7-FeOOH [reactions (4) and (6)] whereas 2z electrons are liberated during oxidation [reaction (8)], the following relation exists between the difference of cathodic and anodic charge, AQ, and the amount of Fe304, nb (Fe304):

AO/_____F_ (i- 2~ zI "nb(Fe~OD nF. nF=

(10)

M. STRAT~NN,K. BOHNENKAMPand H.-J. ENGELL

982

If the sample is reduced at -- 400 mV(SHE) until some magnetite has formed and oxidized again (Fig. 12), AQ can be measured as a function of nb (Fe~O4) which is calculated from the corresponding magnetic data. Figure 13 shows that for different experiments a linear relationship exists between AQ and nb (FeaO4). From the slope of the straight line the extent of oxidation of FeaO4, z, is calculated and a value of z = 0.36 results. Presumably only a surface layer of FeaO~ is oxidized to 7-Fe2Os. From the fact that the degree of re-oxidation of the FeaO 4 is about 35 ~o and from the measured total surface area of the sample it can be calculated that the 7-Fe=Os layer on top of the FeaO4 is about 3 nm thick. For a FeaOJyFe~Oa mixed crystal will be formed the layer will be correspondingly thicker and of changing composition. In any case, the dimensions of the layer are reasonable as the diffusion of the iron ions in the lattice of 7-Fe=Oa should follow a high field mechanism and far greater diffusion distances are very unlikely. The formation of 7-Fe=Oa by oxidation of FeaO4 is a reversible process; at -- 400 mV FeaO4 is reformed. CONCLUSIONS

In the electrochemical reduction of rust the initial reaction is the formation of a Fe s+ intermediate (Fig. 14) which can be assumed to be a reduced surface layer on top of the v-FeOOH crystal. The reduced surface layer can dissolve into the electrolyte within the pores and Fe ~÷ ions diffuse out into the cell electrolyte. This process is very

,~.alF

"

Z

% I

0.10-

,

I,

[Fe 2" ] = 106 tool I1 pH = 6.0

z = extent of reoxidati0n

0f Fe30`

"= AQuP

,,~.

/

/

~

~

f,

F,~{

z=O.O) =-

. /'"

J"="

-

"~,~'~O(z =0.36)

0.05-

0

-

~

0.05

I

i

0.10

0.15 n0(Fe30~) nFe

Fro. 13. g : Measured difference between cathodic and anodic charge, A Q, as a function of the amount of FeaO4 before oxidation, nb(FesO,), in Fez+-fre¢ electrolyte at pH = 6.O. Curves: Calculated values of z~Q after equation (IO) for different values of the extent of reoxidatJon of FeaO=, z.

Phase-transitions in rust layers

983

[Fe Ph3]z" • H" • 2H20

FeZoo.2H O

a- FeOOH • H • e •3Phil" ~

T-FeOOH. H'.e ~

{FeOH OH} :reduced surface layer on II- FeOOH U
Fe304* 2HzO

H z : 0.35 , surface

layer on Fe30,

FIG. 14. Summaryof the observed phase transitions during polarization of a rust layer in 0.2 M NasSO, solution. U* :: - 300 mV (10-6 M Fe6+solution), U* = - 200 mV (2 x 10-a M Fe6+ solution).

much enhanced by acid pH and by complexing agents. The reduction of T-FeOOH stops at this stage if the potential is not below -- 300 mV in a 10 -s M Fe~+-solution and not below -- 200 mV in a 2 x 10 -s M Fe2+-solution, respectively. At lower potentials the formation of FeaO~ starts after some of the Fe s+ intermediate has been formed. The intermediate is partly consumed during this process. Consequently, also Fe 2+ ions from the cell-electrolyte can be incorporated into the y-FeOOH during reduction yielding FesO4 as the reaction product. The Fe 2+ intermediate forming the reduced surface layer can reversibly be oxidized to T-FeOOH at a potential of -- 300 mV or higher whereas the formed FesO4 can only partly be oxidized by formation of a surface layer of T-Fe~Os. Oxidation of F%O4 to T-FeOOH as proposed by Evans is not possible. In a real rust layer which is in contact with iron on one side and with air on the other side reducing and oxidizing conditions change within the layer from one surface to the other. According to Evan's model reducing conditions exist within most of the layer if the rust pores are filled with water but oxidizing conditions prevail in dryed-out layers. The "neutral level" between both conditions moves towards the iron during drying out, towards the atmosphere during wetting of the sample. Thus, portions of the rust were reduced in the one case by the reaction sequence: Fe --> Fe~q+. + 2e (11) 2~-FeOOH -F 2e + 2H+ --> 2 {Fe. O H . OH}

984

M. STRATMANN,K. BOHNENKAMPand H.-J. ENGELL

which at strong reducing conditions, i.e. near to the iron surface, can be followed by: 4y-FeOOH + 2 {Fe. O H . OH} --~ 2 FeaO 4 + 4 H20

02) Fe2q+ + 2y-FeOOH ~ FeaO4 + 2H +. If at the place of the reduction of y-FeOOH the access of oxygen becomes favoured by drying out the pores of the rust the conditions change to re-oxidation of the Fe 2+intermediate and the FeaO4, and the reactions: I 2 { F e . O H . OH} + ~, O~ ~ 2 y - F e O O H + H~O and

(13)

1 2 FeaO 4 + ~, O2 ~ 3 7-F%Os can take place. Within the subsequent cycle y-F%Oa can be reduced again to Fea04;

the electrons necessary for this process again are produced by anodic dissolution of iron: Fe --> Fe a+ + 2e

(14) 4 Y-F%Oa + Fe ~+ + 2e ~ 3 FeaO4. Thus, the Fe 2+ intermediate and Fe~O4 as the reduced, y-FeOOH and y-F%Os as the oxidized components, are the constituents o f the oxidation-reduction cycle within the rust layer which was proposed by U. R. Evans.

Acknowledgements--The authors thank Dr. Schluga, Henkel KGaA, Diisseldorf for the performance

of the BET-measurement and Professor Dr. H. Kudielka, Max-Planck-Institut ffir Eisenforschung, for the analysis of the X-ray data. Financial support from the Deutsche Forschungsgemeinschaft is gratefully acknowledged.

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Phase-transitions in rust layers

985

14. H. SCHWI~rERand H. B6mct, J. electrochem. Soc. 127, 15 (1980). 15. K. BOH~KAMP, Tribune du CEBEDEAU 324, 1 (1970). 16. U. R. EvANs, Trans. Inst. Metal Finish. 37, 1 (1960). 17. U. R. EvANs, Nature, Lond. 206, 980 (1965). 18. U. R. EvAns and C. A. J. TAYLOR,Corros. Sci. 12, 227 (1972). 19. H. OKnDA,Y. Hoso! and H. NAITO, Corrosion 26, 429 (1970). 20. I. MATSOSm~U~and T. UENO, Corros, Sci. 11, 129 (1971). 21. J. Sc~w~'smz, Werkstoffe Korros. 23, 648 (1972). 22. E. Ktn~ZE,Neue Hiitte 19, 295 (1974). 23. I. SUZUKI,N. MASUKO,Y. HISS~tATSU,Corros. Sci. 19, 521 (1979). 24. I. SUZUKX,N. MASUKOand Y. HL~MATSlJ,J. electroehem. Soc. 127, 2210 (1980). 25. J. KaiSER,C. BROWNand R. HEIDERSBACH,Environmental Degration of Engineering, Materials in Aggressive Environments, p. 43. Blacksburg (1981). 26. M. COHENand K. HASmMOTO,J. electrochem. Soc. 121, 42 0974). 27. J. E. HILLER, Werkstoffe Korros. 17, 943 (1966). 28. L. Fonmsmo, Corros. Sci. 20, 1251 (1980). 29. W. M. LATIMER,Oxidation Potentials. New York (1952). 30. T. MISAWA,Corros. ScL 13, 659 (1973). 31. A. L. MACKAY,Reactivity of Solids (ed. J. H. de BOER),pp. 571-583. Amsterdam (1960). 32. A. HICKLmOand D. J. (3. IrEs, Electrochim. Acta 20, 63 (1975). 33. J. D. BEaNm., D. R. Dnsotwr^ and A. L. MACKAY,Nature, Lond. 180, 645 (1957). 34. A. F. WELLS,Structural Inorganic Chemistry, 4th edn., p. 527. Oxford (1975). 35. Ibid., p. 456. 36. M. M. LOHmSNOEL,P. K. RlCm'ER and J. W. SCHULTZE,Ber. Bunsenges. Phys. Chem. 83, 490 (1979). 37. E. B. SAI',rOELL,Colorimetric Determination of Trades of Metals. New York (1950). 38. G. H~c,G, Z. phys. Chem. B29, 95 (1935). 39. Gmelins Handbuch der Anorganischen Chemie (System-Nr. 59, Teil D, Erg. Bd. 2). Weinheim (1959).