Kinetics of anodic oxide-film growth on titanium—II. Neutral and alkaline media

Ekctrochimica Acta, 1971, Vol. 16, PP. 1539 to 1553. Persamon Prw-s. Print& in Northern Ireland

KINETICS OF ANODIC OXIDE-FILM GROWTH ON TITANIUM-I. ACID MEDIA*

I. A. AMMAR and I. KAMAL Department of Chemistry, Faculty of Science, University of Cairo, Gim, U.A.R.

Abstract-Studies on the kinetics of anodic oxidation of Ti are made in H,SO,, HNOI, HCl and H90, solutions of varying concentration. Charging curves are measured at constant cd. Formation rates, and reciprocal capacitance data are reported as functions of the polarizing cd. The exponential relation between ion cd and electric field is also applicable in acid solutions. Values are recorded for the electrolytic parameters A and 23. Potential/log (cd) relations are constructed at constant time of polarization, and at constant Q, where Q is the charge passed. While the slope is a function of cd in the first set, linear relations are obtained at constant Q. These allow the calculation of estimates of the pre-immersion oxide thickness. Measurements at high cds indicate breakdown potentials of about 100 V.

R6sum&Etude des cinetiques de l’oxydation anodique du Ti dans des solutions de H,SO,, HNOI, HCI et H,PO, & diverses concentrations. Les courbes de charge sont Ctablies a densite de courant constante. Les vitesses de formation et les valeurs de l’inverse de la capacitance sont rapport&es, en tant que fonctions de la densite du courant de polarisation. La relation exponentieile entre densite de courant ionique et champ electrique est aussi applicable aux solutions acides. Les valeurs des parametres Blectrolytiques A et B sont indiquees. Les relations potentiel/log(cd) sont repr&ent&s & temps constants de polarisation et aussi a Q constant (Q &ant la charge 6coul6e). Tandis que la pente est fonction de la densite de courant dans le premier MS, on obtient des relations lin&ires B Q constant. Celles-ci permettent une evaluation sommaire de 1’6paisseur pr&mmerg6e de l’oxyde. Des mesures B hautes densites de courant indiquent des potentieIs de rupture d’environ IO0 V.

Zusammenfassung-Die Kinetik der anodischen Oxydation von Ti in HSS04, HNOII, HCl und H,POI wird bei verschiedenen Konzentrationen untersucht. Ladungskurven bei konstanter Stromdichte werden aufgenommen; Wachstumsgeschwindigkeiten, und reziproke Kapazit&en werden als Funktion der Stromdichte hestimmt. Die Exponentialbeziehung zwischen der ionischen Stromdichte und der Feldstlrke gilt such in sauren LBsungen. Parameter A und B werden angegeben.

Werte ftir die elektrolytischen Die Beziehungen Potential/log(Stromdichte) werden f&r

konstante Polarisationsdauer und fur konstante geflossene Ladungsmenge Q aufgestellt. Im ersten Fall ist die Steigung eine Funktion der Stromdichte, w&rend man im zweiten Fall llueare Beziehungen erhllt. Diese gestatten die Berechnungen der Oxydschichtdicke vor dem Eintauchen. Messungen bei hohen Stromdichten ergeben Durchschlagspotentiale von ungef$ihr 100 V.

INTRODUCTION

BECAUSE of the increasing industrial application of Ti, interest in the anodic behaviour of this metal has grown in recent years. Hall and Hackermar- measured the anodic charging curves for Ti in NaCl solutions. Johansen, Adams and Van Rysselberghe2 calculated the formation rates for surface oxide films on Ti and other metals in borate solutions. Theoretical interpretations of these results were given by Van Rysselberghe and Johansen Formation of protective films on Ti by anodic polarization in non- oxidizing acids was studied by Cotton .4 Plating with Pt was described by the same author.s Sterns reported the action of passivating inhibitors which polarize the metal into the passive potential range. Similar observations were made with noble metals.‘+ The effect of CU~+,~--~~ Fe3+,s,11 I- ionslo and HN0310*r2 on the passivation of Ti was investigated. Both C1-,13*x4 and F- ions16 were found to cause breakdown of oxide layers on Ti. The effect of hydrogen pick-up and hydride layers on oxide film formation was studied by Otsuka,la and by Tomashov et a1.l’ Measurements were also

* Manuscript received 29 November 1969. 1539

1540 I. A. AMMAR and I. KAMAL

reported in KOH solutions. 18~1B Photoelectric and anodic polarization studies were made by Rozenfel’d et al. 2o The nature of TiOs films produced by enodic polarization was investigated by a number of authors. Anatase, *la22 mixtures of anatase and rutile,= and also brookite2* were reported.

Although extensive kinetic measurements have been made on the anodic oxidation of some valve metals, eg Ta, Nb and Al,26 yet little attention has been paid to the mode of oxidation of Ti. In the rapidly expanding literature on the anodic behaviour and passivity of Ti, only few studies 1-a have been concerned with the kinetics of oxide film growth. In the present paper we report the formation rates and reciprocal capacities for the oxide film growth on Ti in solutions of mineral acids.

EXPERIMENTAL TECHNIQUE

Spectroscopically pure Ti rods were used. Electrical contact was made with a Cu wire tightly wound in a groove at one end of the Ti rod. The electrode was sealed to glass with a resin over the Ti/Cu junction, and the resistance of this junction was checked before use. In some experiments the electrode surface was abraded with 00 emery paper. However, in most of the work, electrodes were etched in HF solution of composition 4 parts by volume 70% HNO,, 1 part 40 y0 HF and 5 parts H,O. Following this treatment, the electrode was washed with hot distilled H,O, and was then immediately used for measurements. A conventional two-compartment electrolytic cell was used in connexion with a saturated calomel half-cell. De-aeration was done with purified N,. For anodic polarization measurements, the electrode was lowered under polarization into the solution, and its surface was made to touch the tip of the reference electrode probe. Experiments were carried out in H,S04, NHO,, HCl and HsPO, solutions of different normalities, purified by pre-electrolysis. Polarization measurements were made with the galvanostatic technique, and the potential/time relations were measured at constant cd using a calibrated valve volt- meter_ Potentials are here referred to the nhe, and cds were calculated on the basis of apparent surface areas. The cell was thermostated at 30 f 0~1°C.

RESULTS AND DISCUSSION

Measurements were made in O-01, 0.1 and 1.0 N H,SO,, HNOB, HCl and in O-1, 1-O and 5-O N H,PO, solutions. In the low cd range, results of potential/time measurements at constant cd are characterized by three potential regions (i) an instantaneous rise of potential (ii) an intermediate range in which the potential rises gradually and linearly with time (this ends with an inflexion), and (iii) a less steep potential/time region, with no indication of a constant steady potential value. This behaviour is shown in Figs. l-4 for 1.0 N acids. The results indicate passivation even in HCl solutions. The first of the above three potential regions occurs within l-2 min and is not, therefore, included in the figures. Unless otherwise stated, the results reported here are for electrodes etched in HF solutions. Free 0, evolution was not detected by the naked eye. However, when the electrode was viewed with a lens for some time, bubbles were found sticking on the surface, and extremely feeble gas evolution was observed, particularly during the third potential region. The electrode surface re- mained bright throughout the duration of the experiments at low cd (about 100 min).

The amount of electricity Q (in mC/cm2) required to bring the electrode to the end of the second potential region is calculated here as a function of cd and concentrations. In H,SO, solutions, these values of Q do not change regularly with the polarizing cd, i.

Kinetics of anodic oxide-film growth on titanium-I 1541

t, min

FIG. 1. Anodic charging curves for Ti in 1.0 N HISOI. I, 15; II, 22.5; III, 30; IV, 45; V, 60; VI, 75; VII, 85 pA/cmS.

f, min

FIG. 2. Anodic charging curves for Ti in 1 *O N HNO,. I, 7.1; II, 10-7; III, 14.1; IV, 21.7; V, 28.8; VI, 73 pA/cm*.

Instead, they are scattered around a mean value that depends on the acid concentration. Results in HN03 solutions indicate a slight decrease of Q with increase of i in the same acid concentration. Similar observations are made with HCl. The situation is, however, different in H3P04 solutions, and Q increases with increase of i

from (a) 14-4 mC/cm2 (i = 10 ,uA/cm2) to 28-8 mC/cm2 (i = 120 pA/cm2) in O-1 N solutions (b) 25.2 mC/cm2 to 28-8 mC/cm2 in 1.0 N solutions for the same range of cd, and (c) 24.0 mC/cm2 (i = 20 ,uA/cm2) to 49-5 mC/cm2 (i = 75 PA/cm9 in 5.0 N solutions. The mean results are given in Table 1, which shows the general behaviour that Q increases with increase of concentration for the four acids.


I I I I 25 50 75 too

f, min

FIG. 3. Anodic charging curves for Ti in 1-O N HCI. I. 10.8; II, 18.4; III, 23-l; IV, 32.3; V, 77.5 PA/cm*.

t, min

FIG. 4. Anodic charging curves for Ti in 1-O N HsPOd. I, lo; II, 20; III, 40; IV, 80; V, 120 ~Alcm’.

TABLE 1. MEAN VALUES OF Q, TO THIE END OF SECOND POTENTIAL REGION

Solution N

HaSO,, O-01 0.1 1.0

HNO,, 0.01 O-1 1 *o

Range of Mean i Q

pA/cm2 mC/cm*

540 20.6 2.5-25 22.0 15-85 31-o 7-70 21.4 8-80 24.8 7-73 29.3

Sohtion N

HCI, O-01 0.1 1.0

H,PO,, 0.1 1.0 5-o

Range of i

pA/cm”

7-66 6-49

1 O-80 lo-120 lo-120 20-75

Mean Q

mC/cm*

17-S 20.6 26-5 21-1 27.3 36.3

Kinetics of anodic oxide-film growth on titanium--I 1543

Upon switching off the current during the second potential region, the anodic polarization decays within 10 min, and the potential remains constant on open circuit. Repolarization at the same cd produces a sudden jump (2 min), and the potential/ time relation obtained is a continuation of that before anodic decay. Hall and Hacker- man1 attribute this phenomenon to the irreversible chemisorption of neutral 0 atoms to form 0.85 of a monolayer, after the charging of the double layer is completed. If chemisorption of oxygen is the main cause of passivity, spontaneous passivation accompanied by the shift of potential of non-polarized Ti in the positive direction would be expected to occur more readily in aerated than deaerated solutions. Figure 5

0

0

s! s I;; -0.

-O-

-@a I I I I I

30 60 90 120 I50

f. min

FIG. 5. Open-circuit E,/r relation in 1-O N acids. I, HCl; II, H&O&; III, HsPOd; IV, HNO*.

shows the results of open-circuit potential/time measurements in thoroughly de-aerated 1.0 N solutions. For both H$O, and HCI, the potential decreases with time and reaches the steady state values -4480 mV and -560 mV for the two acids respectively. However, the opposite trend is observed in the results of H,PO, and HNO, solutions. Here, the potential increases with time to a steady value of -30 mV for HsPO, and +410 mV for HNO,. The latter behaviour cannot be attributed to chemisorption of oxygen, but rather to the oxidation of Ti surface. It is likely that chemisorption of oxygen is rapid and occurs during the first instantaneous potential rise, prior to the anodic oxide film formation. Hackermanz6 considers adsorption as a primary process that leads to the formation of an oxide film whose nature (crystalline or amorphous) determines the subsequent passive or active behaviour of the metal.


The steady state potential of Ti is affected not only by the electrochemical proper- ties of the metal, but also by the nature of the surface oxide film formed under the prevailing experimental conditions. The stability of this film depends on the type of electrolyte, its concentration, and temperature. Self-passivation has been reported for Ti in dilute H$O, and HCl solutions. IQ In this case, oxide-tim formation occurs due to oxidation of Ti, and the rate of this process exceeds the oxide dissolution rate. However, in concentrated H,SO, and HCl solutions, self-passivation does not occur as a result of the increased solubility of oxide.ls

Before Ti is first dipped in solution, it is likely that a pre-immersion oxide layer of thickness 8, occurs on its surface. This is probably the case for most procedures of surface pre-treatment in which the metal is left in contact with air for some time. Provided that the anodic process is entirely the formation of an oxide film, anodic polarization at constant i leads to film growth according to

A8 = (1)

where M is the molecular weight of oxide, n the number of faradays Frequired for the formation of one mole oxide, s its density, G the roughness factor, i the time, and r the volume of oxide formed per C. The latter is sometimes known as the anodic oxidation electrochemical equivalent_ Film growth may occur by the transport of positive ions through oxide to the oxide/solution interface, or by transport of oxide ions in the opposite direction. Since the time is reckoned from the moment the polarization is started, the oxide thickness at any given time t is expressed by

6 = 6, + ; (2)

Oxides of Ti, particularly the dioxide, are poor electron conductors, and charge transport is predominantly ionic. l This is generally true except at very low electric fields, and at breakdown, where the contribution of electron conductivity is significant. During film growth, the ion cd is an exponential function of the field strength n,27

i = A exp (BX?), (3)

where A and B are the electrolytic parameters. The behaviour of a practical system approaches this exponential relation when the ion flux in the direction opposing the effect of field is negligible.

Dependent on the location of the energy barrier for ion transport, there are essen- tially two models (i) the Cabrera-Mott modelas in which the energy barrier occurs at the metal/oxide interface, and (ii) the Verwey mode129 associated with a barrier within the oxide. Experimental evidence in support of either model is found in the literature.3o Simultaneous control by the two energy barriers has been proposed by Dewald31 over certain ranges of temperature and field. Vermilyea et aZ.32*53 report convincing evidence that (i) the rate-determining step in the formation of Ta,O, is within the film alone (ii) the mobile ions are interstitial, and (iii) the concentration of these ions is dependent on field.

It appears that the behaviour of Ti is in some respects analogous to that of Ta, particularly in connexion with the occurrence of a positive space-charge within the

Kinetics of anodic oxide-film growth on titanium-1 1545

0xide.l The contribution of the two interfacial potentials at metal/oxide and oxide/ solution boundaries to the total pd E (measured against a suitable reference) may be considered constant (at constant i) at least in the range of film formation. Hence,

E = Et + + = Ei + const, (4)

where Ei is the pd across the film, and + is the sum of interfacial contributions. The increase of potential is, therefore, manifested across the oxide layer such that

W = E,/d. Hence, from (2), (3) and (5)

i = A exp

The dependence of E on i and t follows from (4) and (6)

(5)

(6)

(7)

When the oxide thickness increases with time at constant i, Ef and consequently E must increase in order to maintain the field and the current. Hence,

This represents the formation rate. Except in very few cases, the linearity between E and t is experimentally verified for Ti in acid solutions in the intermediate potential region (Figs. l-4).

Formation rates are calculated from the expanded E/t relations before 0, evolution. The logarithm of formation rate is linearly related to log i for all acids and concentrations studied (Figs. 6-9). For uniformity, the formation rate is expressed in V/s and i in A/cm 2. The slope of this linear relation is l-18 for H2SO4, 1.26 for HNO, and 1-16 for HCl, independent of concentration. However, in HsPO, solutions the slope varies slightly with the acid concentration being 1.2 in 0.1 and 1.0 N and 1.34 in 5 N solutions. In all these results, increase of acid concentration decreases the intercept of the straight line. In Fig. 6, results on abraded electrodes in 0.1 N H,SO, are included for comparison. In this case, a straight line is also obtained which is nearly parallel to the other results. However, its intercept is smaller than the corresponding values obtained with electrodes etched in HF solutions. The linearity between log (dE/di) and log i suggests the empirical relation

(dE/dt) = a(i)“,

with a and b as constants. A similar relation was given by Vermilyea34 for the dependence of (ddjdt) on i for Ta. According to (8), this linearity is not expected to hold over a large range of cd. However, in a limited cd range, the difference in the log [2-303 log (i/A)] term is relatively small, and hence the lines are linear, as experimentally observed.

1546 I. A. &MAR and I. KAMAL

-log i, A /cm*

FIG. 6. Relation between log (formation rate) and log (cd) for HISO~. I, 0.01; II, O-1; III, 1-O N; IV, O-1 N, abraded electrode.

zo-

c > 25-

,_

$

$ +

40 .

- log i. A/cm’

FIG. 7. Relation between log (formation rate) and log (cd) for HNO.. I,O*Ol; II,O-1; III, l-ON.

The unitary formation rate (Ri) or reciprocal capacitance (l/C) is obtained by dividing both sides of (8) by i,

1

c (9)

where C is the capacitance unit area. If the reciprocal capacitance is expressed in cm2/F, ie Ri in V. cmZ/A.s, and i in A/cm 2, the linear relation between (l/C) and log i gives a slope [2-303 r/(aB)] and an intercept on the reciprocal capacitance axis of


- logi, A/cm2

FIG. 8. Relation between Iog (formation rate) and log (cd) for HCI. I, 0.01; II, 0.1; III, 1-O N.

- 2.5- ‘2

P 4r E

H I

l I I I 459 4.4 40

- logi, A /cm’

FIG. 9. Relation between log formation rate and Iog (cd) for H,PO,. I, O-1 ; II, 1.0; III, 5.0 N.

- [2.303 r/(oB)] log A. The results are shown in Figs. 10-13, with C in ,uF. Taking M as 79-9 g/mole for TiOz, and s = 3.84 g/cm3 for anatase, the value of t becomes 5-4 x 10- s cm3/C. With this value of r, and assuming (T = 3, the electrolytic parameters A and B are calculated from the lines of Figs. 10-13. The results are given in Table 2, which shows that (i) A increases with increase of concentration for the four acids (ii) B is practically independent of concentration for H,SO*, HCl and &PO,, and (iii) a slight increase of B is observed with increase of concentration for HNO,, as is evident from Fig. 11. The values obtained with the abraded electrode in O-1 N H,SO, are


1 I I I

50 4.5 40

-log i, A /cm2

FIG. 10. Reciprocal capacitance us log (cd) for HISOp. I, O-01 ; II, O-1 ; III, 1-O N; IV O-l N, abraded electrode.

o.0751 50 4-5

---logi. A/cm2

FIG. 11. Reciprocal capacitance us log (cd) for HNO,. I, O-01 ; II, O-1 ; III, 1.0 N.

A = l-7 x lo-8 A/cm2, B = l-4 x lo-8 cm/V. Allowing for the difference in units, our (l/C) values are of the same order of magnitude as those reported by Johansen,

2 Adams and Van Rysselberghe for Ti in ammonium borate solutions (pH = 8). Agreement is also observed with B. However, our values of A are considerably higher than those obtained in borate solutions. This difference may be related to the pH of solution, since this parameter decreases with increase of pH (Table 2). The relation between field strength H and (l/C) follows from (3) and (9),

0 1 J-g= - 0 r i5’

(10)

and U is directly calculated from reciprocal capacitance data. The values obtained are


O-20 -

‘: 2

m x

9 r.ii

5 015-

% -

0.10 -

L I I I 5.0 45 4-o

-logi, A /cm2 FIG. 12. Reciprocal capacitance us log (cd) for HCI.

I, 0.01; II, 0.1; III, 1.0 N.

-log i, A/cm2

FIG. 13. Reciprocal capacitance us Iog (cd) for H,PO,. I, O-l ; II, 1.0; III, 5.0 N.

of the order 1 O6 V/cm. The activation distance or half-barrier width Q* is calculated from the values of B in Table 2 according to B = +z*/RT,~” where q is the charge on the ion. The results obtained lie between O-28 and 0.54 A. The corresponding values reported by Van Rysselberghe and Johansen from data of aB lie between 3.55 and 4.35 A.

The E/log i relations during oxide film growth are deduced from (7) if the experimental conditions of pre-treatment lead to a constant pre-immersion oxide thickness S,,. Hence, at constant t,

(g),= (2) +i($)(l +lni), (11)


TABLE 2. ELECTROLYTIC PARAMETERS

Solution A B Solution A B N A/cm” m/V N A/cm8 cm/V

H,SOI, O-01 5.7 x 10-S o-1 Q-7 x 10-B 1-o 3.4 x 10-7

HNO*, 0.01 s-3 x lo-’ o-1 5-7 x 10-1 I.0 5.8 x lo-’

7.2 x 1O-7 6-9 x 10-1 7.2 x lo-’ 4.3 x lo-’ 5.0 x lo-' 6-l x lo-'

HCI, 0.01 0.1 1-o

HgPOI, O-1

::;

2-3 x lo-* 6.9 x lo-' 2-4 x lo-* 7-4 x 10-1 1-l x lo-’ 6.9 x lo-' 7-6 x lo-’ 7-s x 10-1 2.5 x 10-T 7.2 x lo-’ 2.5 x 10-e 8-3 x IO-’

and the slope is a function of i, thus increasing with increase of cd. This behaviour is

shown in Fig. 14. Therefore, the E/log i relations are calculated at constant Q. For this purpose, E/Q relations are first constructed at constant i [cf(9)]. An example of

I I I I I 5-2 5.0 4.8 4.6

-log i, A /cm*

FIG. 14. Potential/log cd relations at constant f. a, 5 min; 0, 7-S min, x , 10 min.

I, 1.0 N HNOs II, O-01 N HNOa.

these relations is shown in Fig. 15 for Ti in O-01 N HCI. The E/log i relations are then calculated from the results of Fig. 15 at several constant values of Q. Straight lines are obtained in the low cd range (Fig. 16), and deviations from linearity occur at higher cds. The slope of this linear relation is [cf (7) with Q = it]

= 2.303[@ + (g], (12)

and it increases with increase of Q. Equation (12) is used together with Fig. 16 in

Kinetics of anodic oxide-tihn growth on titanium-I 1551

0 I I I 5 IO 15

0. mC/cm*

FIG. 15. Potential/Q relations at constant cd in a01 N HCl. I, 6.8; II, 11-2; III. 15; IV, 22.1; V, 30.3; VI 66.2 pA/cm’.

-leg f, A/cm’

FIG. 16. Potential/log (cd) reIations at constant Q in DO1 N HCl. I, 8; II, 12; III, 16 m c/ems.

order to obtain estimates for the pre-immersion oxide thickness 6,. The results are relatively low and lie between 4 and 5 A. These are then used in (7) in order to calculate the sum C$ of the two interfacial pds. The resulting values lie between -400 and -600 mV. Similar calculations have been made with other acids, and slightly higher values are sometimes observed for 6,.

1552 I. A. AMhWR and I. KAMAL

No direct means is, however, available to check the accuracy of the above values, in particular the sign of 4. The situation is not simple, since all calculations are made in the intermediate potential region following the rapid potential rise. Experimental evidence shows that the phenomena occurring in the two regions are not the same. Therefore, extrapolation of the data calculated in one region to the other can only be considered as indicative. Moreover, three other sources of error may be added (i) the assumption of constant 4 (ii) the assumption of a 100 per cent efficiency for film formation (iii) the neglect of any contribution from surface hydride tims formed during electrode pre-treatment. 16p1’ Such hydride films not only modify the kinetics of oxide- film formation, but may also lead to delay of O2 evolution. This is indeed the case, since very feeble 0s evolution is observed at potentials much higher than the thermo- dynamic values (Figs. l-4). Also, O2 evolution occurs more readily on abraded than on etched electrodes. Although Ti is not an ideal valve metal, yet the present treatment shows that the kinetics of oxide-film formation on Ti resemble those on more nearly ideal valve metals like Ta. In the treatment proposed by Van Rysselberghe and Johansen, the authors use the Cabrera-Mott model and neglect the possibility of any space-charge within the oxide. They suggest a constant field from the bulk of solution to the bulk of metal and consider the rate-determining step to be either (i) transfer of metal ions from metal to oxide (at intermediate cds), or (ii) transfer of oxide or OH- ions from solution to oxide (at very low cds). On the other hand, our results are best explained on the basis of a net space-charge in the oxide phase.

Measurements at high cds (mA range) indicate three regions (i) a linear rapid rise of potential which ends with a slight maximum (ii) evolution of O2 at this maximum ;

the potential then decreases gradually, remains constant for some time, and eventually starts to increase gradually again (iii) in the third region, the potential rises rather rapidly, reaching 100 V before oxide breakdown begins. Interference colours are observed, and the electrode surface is dark grey at the end of the experiment. The potential at the maximum increases with increase of i. The difference between the breakdown observed here (100 V) and that found by Hall and Hackermanl (10 V) is attributed to the difference in the purity of Ti, the procedure of pre-treatment, and the nature of electrolyte. Hall and Hackerman explain the nearly constant potential region as due to the formation of TiOs together with O2 evolution. However, the present results indicate that the formation of TiO, occurs before O2 evolution. The electron-current/ion-current ratio is not constant throughout the range of polarization. It is possible that the ion-current predominates during oxide film formation. As the thickness of oxide increases, crystallization may start, and this increases the defect character of oxide.35 Therefore, the electron-current increases and 0s evolution occurs on grain boundaries and lattice imperfections. This leads to chemical oxidation at pores, and when the crystalline oxide becomes complete the electron-current decreases again and 0s evolution ceases. At this stage, electrical conduction through oxide is mainly ionic, and the thickness starts to increase again until breakdown finally occurs.

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Kinetics of anodic oxide-film growth on titanium—II. Neutral and alkaline media

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Transcript of Kinetics of anodic oxide-film growth on titanium—II. Neutral and alkaline media