THE ANODIC DISSOLUTION OF MAGNESIUM IN …library.nmlindia.org/FullText/CS3997101981.pdf · THE ANODIC DISSOLUTION OF MAGNESIUM IN CHLORIDE ... The next step in the study of the corrosion

Corrosion Science, Vol. 39, No. l&l 1, pp. 1981-2004, 1997 0 1997 Elsevier Science Ltd

Printed in Great Britain. All rights reserved 001&938X/97 $17.00+0.00

PII: soolo-938X(97)ooo9o-5

THE ANODIC DISSOLUTION OF MAGNESIUM IN CHLORIDE AND SULPHATE SOLUTIONS

G. SONG,*+ A. ATRENS,* D. ST JOHN,* X. WUt and J. NAIRN*

* CRC for Alloy and Solidification Technology (CAST), Department of Mining, Minerals and Materials Engineering, The University of Queensland, Brisbane, Qld 4072 Australia

+ The State Key Laboratory for Corrosion and Protection, Institute of Corrosion and Protection of Metals, Chinese Academy of Sciences, 62 Wencui Rd, Shenyang, 110015, China

Abstract-The electrochemical behaviour of magnesium was studied in representative chloride and sulphate solutions including NaCI, Na2S04, NaOH and their mixed solutions, HCI, and H2S04: (1) by measuring electrochemical polarisation curves, (2) by using electrochemical impedance spectroscopy (EIS), and (3) by simultaneous measurement of hydrogen gas evolution and measurement of magnesium dissolution rates using inductively coupled plasma atomic emission spectrophotometry (ICPEAS). These experiments showed that a partially protective surface film played an important role in the dissolution of magnesium in chloride and sulphate solutions. Furthermore, the experimental data were consistent with the involvement of the intermediate species Mg+ in magnesium dissolution at film imperfections or on a film-free surface. At such sites, magnesium first oxidised electrochemically to the intermediate species Mg+, and then the intermediate species chemically reacted with water to produce hydrogen and Mg*+ The presence of Cl- ions increased the film free area, and accelerated the electrochemical reaction rate from magnesium metal to Mg+ 0 1997 Elsevier Science Ltd

Keywords: A. magnesium, B. EIS, corrosion, electrochemical dissolution, negative difference effect.

INTRODUCTION

The use of magnesium alloys has been accelerating in automotive applications where advantage can be taken of the low density of magnesium. However, because magnesium is a very reactive metal, corrosion resistance continues to be an important issue. Consequently, it is important to understand the anodic dissolution mechanism for the corrosion of magnesium and its alloys.

The influence of environmental factors on magnesium corrosion has been intensively investigated.lm5 For example, it is well known that Cl- promotes the corrosion of magnesium in aqueous solutions,3’4 and that the corrosion rate usually increases with increasing Cl- concentration.’ SO4’- is believed to have much less influence than C1-,6 and

so4 2- is sometimes used as a supporting electrolyte for electrochemical investigations.’ However, most of these prior studies only measured the corrosion rate and did not consider the actual anodic dissolution mechanism. Consequently, there is no clear understanding of the role of Cl- in the mechanism of magnesium corrosion, even though Cl- is one of the factors of most concern to magnesium alloy users. The rapid corrosion of magnesium in Cl- containing solutions is often simply attributed to an accelerated breakdown of a surface film on magnesium by Cl-.

Manuscript received 19 March 1997; accepted 27 May 1997.

1981

1982 G. Song et al

It is clear that Cl- accelerated film breakdown is part of the overall anodic dissolution process in neutral and alkaline solutions. However, it is also to be expected that Cl- ions have an influence on the reaction sequence by which metallic magnesium is converted to

Mg 2 T ions during the corrosion process. The following four models have been proposed to describe the overall magnesium

dissolution process: (i) partially protective film mode1;1~*~‘2 (ii) hydride mode1;8~‘5-18 (iii) univalent magnesium ion model ‘,19 22 and (iv) undermining second phase particle modeI.8’2’ 25 Our previous study I2 of the electrochemical corrosion of pure magnesium in 1 N NaCl found that the corrosion behaviour was consistent with a partially protective surface film. The surface film covered all of the surface for applied potentials more negative than a critical potential Ept which was about 15 mV more negative than the corrosion potential E,,,,. Above Ept the surface film was only partially protective and its surface coverage decreased with increasing applied potential, E. This prior study l2 provided strong evidence that a partially protective film played a key role in the corrosion process of pure magnesium in NaCl solutions of pH 3-l 1.

The next step in the study of the corrosion behaviour of pure magnesium is to understand the detailed reactions that occur at the locations where magnesium is directly exposed to corrosive solutions, either at the imperfections in the protective film or at the locations where there is no surface film. That is the purpose of the present study which has employed three different complementary electrochemical techniques to obtain a clear understanding.

Potentiodynamic polarization curves were measured in a range of representative chloride and sulphate solutions in order to gain an overview of magnesium corrosion in these solutions.

Electrochemical impedance spectroscopy (EIS) measurements were carried out in two solutions at carefully selected potentials in order to study the detailed reaction sequence of magnesium corrosion.

Gas collection and ICPAES experiments were carried out in a range of representative solutions. The gas collection and ICPAES experiments allow simultaneous measurement of the rates of both the anodic and cathodic reactions and therefore were very useful in further elucidating the reaction mechanism, and in particular the phenomena of the negative difference effect.

The negative difference effect is considered in some detail in the prior work i2, which includes a detailed description which is not repeated herein. The essential feature of the NDE is that the rates of both the anodic and cathodic reactions increase with an increase of applied potential. The NDE is very different to ‘normal’ electrochemical behaviour, as for example an iron specimen in acidic solutions where an increase in potential causes an increase in the rate of the anodic reaction but causes a decrease in the rate of the cathodic reaction. (NDE).

The discussion section relates all this experimental data to the proposed reaction sequence for magnesium corrosion.

EXPERIMENTAL METHOD

The material used for this study was the same as that used in the prior study:12 99.96% pure magnesium ingots whose main impurities as detected by inductively coupled plasma

Magnesium chloride and sulphate solutions 1983

atomic emission spectrophotometry (ICPAES) were: 240 wppm Fe, 140 wppm Al, < 10 wppm Ni, < 10 wppm Cu, 10 wppm Mn, 20 wppm Ca, 30 wppm P, and 10 wppm Ba. Specimens were cubes 10 x 10 x 10 mm. Each cube was moulded into epoxy resin with only one side unsealed, so that the working area exposed to the electrolyte was 10 x IO mm. Before use, the electrodes were ground with 600 grit Sic paper, washed with distilled water and dried by warm flowing air.

The following solutions were used: (i) 1 N NaCl, pH = 11; (ii) 1 N Na2S04, pH = 11; (iii) 1 N NaCl + xN Na2S04, or 1 N Na2S04 + xN NaCl, pH = ll,O~x~0.05; (iv) 1 N NaCl, pH = 13; (v) 1 N Na2S04, pH = 13; (vi) 1 N NaOH; (vii) 1 N NaOH + zN NaCl or 1 N NaOH + zN Na2S04 01~510; (viii) 1 N HCl; (ix)1 N H$Ob. All solutions were made with A.R. chemicals and distilled water. The pH was adjusted to the desired value with 1 N NaOH. All solutions were aerated.

Potentiodynamic polarisation curves were measured in a glass cell with the pretreated magnesium specimen as the working electrode. The counter electrode was a platinum gauze (25 mm x 25 mm, 52 mesh) and a saturated calomel electrode (SCE) was used as the reference electrode. The potential was scanned from - 1870 mV at a rate of 0.2 mV/s using a potentiostat, Potentio-Galvano-Scan Wenking PCS 81. The corresponding current passing through the electrode was collected from the potentiostat by a high resolution data logger, PICO ADC- 16.

Electrochemical impedance spectroscopy (EIS) of magnesium was measured in solutions (i) 1 N NaCl, pH = 11 and (ii) 1 N Na#Od, pH = 11. First the magnesium electrode was polarised at a constant potential in the test solution for 5 min. Then, EIS measurements were carried out using a Solartron-Schlumberger 1286 frequency response analyser. The AC amplitude of perturbing signal was 5 mV, and the frequency was varied from 19 kHz to 10 mHz.

Gas collection and inductively coupled plasma atomic emission spectrophotometry (ICPAES) experiments were conducted to measure the amount of evolved hydrogen gas and the amount of dissolved magnesium from the magnesium electrode. The sample was put into the cell which contained about 450 ml of test solution, and polarized at a constant potential or current density. The corresponding current density or electrode potential was recorded as a function of time. At the same time, a funnel over the electrode led all gas bubbles from the specimen into a burette which was vertically mounted over the funnel. The burette was initially full of the test solution. The entry of the gas displaced the test solution in the burette and so the evolved gas volume could be measured directly by reading the position of the liquid level. At the end of the experiment, the specimen was removed from the cell. The solution left in the burette was poured back into the cell. Approximately 20 ml of 98% H2S04 was added to the electrolyte in order to dissolve the white corrosion product. The electrolyte was made up to 1 litre of solution with distilled water and this solution was analysed by ICPAES to determine the concentrations of all species which had dissolved from the magnesium specimen. A Spectra ICPAES model M + P instrument was used for the ICPAES analysis.

All the above experiments were carried out at a constant temperature 30 + 1 “C, and all the potentials were referred to the saturated calomel electrode (SCE).

1984

Polarisation curves

G. Song et al.

EXPERIMENTAL RESULTS

Figure 1 illustrates typical polarisation curves (A and B) for magnesium in the two solutions: 1 N NaCl and 1 N Na2S04, both at pH 11. These polarisation curves did not exhibit Tafel behaviour in the anodic branch. In contrast, the cathodic branch exhibited linear Tafel behaviour for potentials more negative than a critical potential Ept. Some small pits were observed to initiate at the critical potential and at potentials more positive than the critical potential, but there were no pits at more negative potentials. It was also observed that hydrogen evolution occurred mainly at these pits. Curve A for magnesium in 1 N NaCl was generally higher than curve B for magnesium in Na2S04, and moreover, the slope of the anodic branch was much steeper than the slope of the cathodic branch. The corrosion potential (EC,,,) of magnesium in 1 N NaCl was more negative than E,,,, in 1 N Na2S04.

Polarisation curves were measured in eight different solutions of the general type 1 N NaCl + xN Na2S04, or 1 N Na$Od + xN NaCl, pH = 11, 0~~~0.05. In each case the electrode surface was severely corroded and uneven after the experiment. All these polarisation curves had shapes similar to the two typical curves A and B, were located between these two curves A and B, and there was a systematic shift of the polarisation curves from curve A to curve B. Increasing the NaCl concentration in the 1 N Na2S04 solution increased both the anodic and the cathodic current densities of magnesium, whereas increasing the Na$Od concentration in the 1 N NaCl solution decreased both the anodic and cathodic current densities. In all cases, EC,,, was always only slightly more positive than Ept and both showed similar trends with solution concentration, as illustrated in Fig. 2. Increasing NaCl concentrations caused an increase in both Ecorr and Ept, whereas an increase in Na2S04 concentration caused a decrease in both EC,,, and Ept. The maximum difference between EC,,, and Ept was 17 mV, and Ept was located within the linear region around EC,,,. The polarisation resistance, R,, at Ecorr was measured by linear regression. R,

decreased with increasing NaCl concentration, and increased by a small amount with increasing Na2S04 concentration (see Fig. 3).

A 1N NaCI. pH=ll B IN Na,SO,, pH=ll C 1N NaCI. pki=13 D 1N Na,SO,. &+I=13

10-3 I I -1800 -1700 -1600 -1500 -1400 -1300

Applied Potential, E [mV]

Fig. I. Typical polarisation curves (A. B, C and D) for magnesium in the following solutions, (A)

1 NNaCI.pH=II.(B)l NNa2S04,pH=ll,(C)1 NNaCI,pH=13,and(D)l NNa2S0,,,pH=13.


-1540

-1560

-1580 s k W -1600 2 $ f -1620

z 0.

-1640

-1660

Econ (mV. 1N Na,SO, + xN NaCI)

I

0 I

0.01

I I I

0.02 0.03 0.Ci-l 0.05

Solution Concentration, x ( N)

Fig. 2. The dependence of the corrosion potential (E,,,,) and the pitting potential (E,,) on NaCl concentration in 1 N Na$X$ and on Na$SOd in 1 NNaCl at pH 11. _-

350

300

’ 5 250

$ 200

g 150 v1 E 100 g

50

0

bc 1 N N&SO, + xN NaCl

Rp- 1 N N&SO, + xN NaCl

0 0.01 0.02 0.03 0.04 0.05

Solution Concentration x (N)

Fig. 3. The dependence of the linear polarisation resistance (R,_) and the cathodic Tafel slope (b,) on NaCI concentration in 1 N Na2S04 and on Na2S04 in 1 N NaCl at pH 11.

The cathodic Tafel slope, bpc, was obtained from the Tafel region of the cathodic branch of the magnesium polarisation curves. Figure 3 shows that there was only a slight decrease of 6,, with an increase of Na2S04 concentration. The b, values varied from about 260 mV/ dec in 1 N Na2S04 to about 315 mV/dec in 1 N NaCl. These values were close to the cathodic Tafel slope value of 280 mV/dec obtained by Penere et ~1.’ for magnesium in 1 N Na2S04. They were much higher than the values measured in typical electrode systems such as iron in acidic solutions where values about 120 mV/dec were usually obtained.28-30 This indicates that the mechanism of the hydrogen evolution reaction on the surface film on magnesium could be significantly different to the mechanism on a bare metal surface as for iron in acidic solutions.

1986 G. Song er al.

The cathodic process within the Tafel region was mainly hydrogen evolution.‘* In a

static solution without stirring, the current associated with reduction of oxygen can be neglected compared with that due to hydrogen evolution as it is controlled by the diffusion of molecular oxygen in the solution. The exchange current density, I,,, of

hydrogen evolution on the film covered magnesium was deduced by extrapolating the cathodic Tafel line to -405.2 mVsce which is the equilibrium potential of hydrogen

evolution in a pH 11 solution. Figure 4 shows that ZPc decreased significantly with an increase of Na2S04 concentration and increased only slightly with an increase in NaCl

concentration. Similar polarisation curves were also measured in 1 N NaCl and 1 N Na2S04, both at

pH 13; these are shown as curves C and D in Fig. 1. After polarisation, localised corrosion

damage was also observed on these electrodes. Comparison of curves A and C showed that the current density of the anodic branch of the polarisation curve in 1 N NaCl at pH 11 was

much higher than in 1 N NaCl at pH 13. The increase in the cathodic curve at pH 13 going from 1 N NaILSO to 1 N NaCl was small compared with the larger increase for the pH 11

solutions. Figure 5 illustrates the polarisation curves in the more alkaline solutions: 1 N NaOH,

1 N NaOH + 2 N NaCl, and 1 N NaOH saturated with NagSOd. The cathodic branches were nearly the same in all these three solutions, whereas there were significant differences in

the passivation behaviour in the anodic parts of these polarisation curves. The passivation region was wide in 1 N NaOH and in 1 N NaOH saturated with Na,SO,; above about

2000 mVsce, the anodic current density increased sharply, accompanied by oxygen evolution

0.01 0.02 0.03

Solution Concentration, x (N)

1 N NqSO4 + xN NaCl

Fig. 4. The dependence of the exchange current density (I,) of the cathodic hydrogen evolution

reaction on NaCl concentration in 1 N Na2S04 and on Na2S04 in 1 N NaCI at pH 1 I,


Potentlal, E [mV]

Fig. 5. Polarisation curves for magnesium 1 N NaOH, 1 N NaOH + 2 N NaCl, and 1 N NaOH saturated with Na2S04.

from the electrode surface. The electrode surface was almost unchanged after anodic polarisation in these two solutions. In contrast, the passivation region of magnesium was narrow in 1 N NaOH + 2 N NaCl. At about - 750 mV, the anodic current density suddenly increased, and pits appeared on the surface of the magnesium electrode. A large amount of gas evolved from these pits, and the pit size and number increased significant with increasing potential. Figure 6 illustrates the dependence of corrosion potential of magnesium on NaCl and Na2S04 concentrations in 1 N NaOH. E,,,, decreased with

-1200

5 ‘= -1300 t:

3 -1400 . 1 3 -1500

% .g -1600 8 t S

1 N NaOH + zN Na$O,

1 N NaOH + zN NaCl

I i 1 I 1 I I 1 I

012345678 9 10

Concentration (N) Fig. 6. The dependence of the corrosion potential (E,,,,) of magnesium on NaCI and Na2SG4

contents in I N NaOH.

I988 G. Song et al.

-2800 -2600 -2400 -2200 -2000 -1800 -1600 -1400 -1200

Applied Potential, E [mV] Fig. 7. Polarisation curves for magnesium in 1 N HCI and in I N HLS04.

increasing NaCl concentration, whereas E,,,, increased very little with increasing Na2S04

concentration. Figure 7 presents the polarisation curves for magnesium in 1 N HCl and in 1 N H2S04 in

a semi-logarithmic presentation. They were almost straight lines in a very wide range from

- 2500 mVsce to - 1300 mVsce if displayed with linear axes. The current densities in these acid media were significantly higher than in the solutions with higher pH values. Furthermore, the cathodic current densities in 1 N HCl were much higher than in 1 N

H2S04. After each polarisation experiment, the electrode surface was bright and uniform, even though a significant amount of electrode had been consumed by corrosion.

Electrochemical impedance spectroscopy Figures 8 and 9 present Nyquist plots for magnesium in 1 N NaCl and 1 N Na2S04,

respectively, at different potentials. At a cathodic potential, E = EC,,,-90 mV, there was

-2i (Ohm) -Zi (Ohm) -ZI (Ohm)

E=Econ-WmV 103

Zr (Ohm) Zr (Ohm)

Fig. 8. Nyquist plots for magnesium in 1 N NaCl at different applied potentials, (a) E,,,,-90 mV,

(b) Lx, and (c) E,,,, + 90 mV.


-21 (ohm) -21 (ohm)

~~~~~~~~~~~

Zr (ohm) Zr (ohm) Zr (ohm)

Fig. 9. Nyquist plots for magnesium in 1 N Na2S04 at different applied potentials, (a) E corr - 90 mV, (b) IL,, and (c) E,,,, + 90 mv.

only one capacitive loop for magnesium in both 1 N NaCl and 1 N Na2S04. At the corrosion potential there were two obvious capacitive loops in the high and low frequency regions in both solutions. At an anodic potential, E= EC,,, + 90 mV, there was a capacitive loop in both solutions in the high frequency range but different behaviour in the middle and low frequency ranges. In 1 N NaCl, there was an inductive loop in the middle frequency range, and a capacitive loop in the low frequency region. In contrast, in 1 N Na2S04 there was a capacitive loop in the middle frequency region, and some scattered data in the low frequency region, indicating that there was an inductive loop in the low frequency range. These Nyquist plots for magnesium in 1 N Na2S04 were similar to the published results 7 which showed a clear inductive loop in the low frequency range.

The capacitive loop in the high frequency region is always related to the transient resistance (R,) and the double layer capacitance (Cdl) of the electrode.73’3v’4 The Cdl and R, were calculated from the Nyquist plots and their values are shown in Figs 10 and 11. Cdl increased with increasing electrode potential from cathodic to anodic. R, exhibited a maximum at around E,,,,, with lower values at both cathodic and anodic potentials.

80

70

60

50

-100 -50 0 50 100 E - Ecorr (mV)

Fig. 10. The dependence of transient resistance (&) and double layer capacitance (C,,) of the magnesium interface on electrode potential (E) in I N NaCI.

G. Song et al.

-50 0 50 100 E - Ecorr (mV)

Fig. Il. The dependence of transient resistance (R,) and double layer capacitance (C,,) of

magnesium interface on electrode potential (E) in 1 N Na2S04.

Gas collection and ICPAES experiments Figure 12 presents the data for hydrogen evolution rates and magnesium dissolution

rates for magnesium at different applied current densities in 1 N NaCl and 1 N Na2S04. In both solutions, the hydrogen evolution rates increased with increasing applied

current density indicating that there was a negative difference effect (NDE).t2 At the corrosion potential and for an applied anodic current density, the hydrogen evolution rates and magnesium dissolution rates in 1 N NaCl were higher than those in 1 N

Na2S04. Under the condition of an applied cathodic current density, the magnesium dissolution rates were almost zero and the hydrogen evolution rates were almost the

same in both solutions.

200 ,

180 +

E 160

-z 2 140 3

t $% g

120

100

=- ;a 80

.o r 60

-Cc 2 40

:: .Y 20

0 0 i

q gas evolution m Na,SO,

n Mg dissolution in NaCl

q gas evolution in NaCl

t -1 0 30

Applied Current Density (mAlcm2)

Fig. 12. Hydrogen evolution rates and anodic magnesium dissolution rates at different applied

current densities in 1 N Na$04 and in I N NaCl at pH I I


Figure 13 shows that both the hydrogen evolution rate and the magnesium dissolution rate increased with increasing applied anodic current density in 1 N HCI and 1 N H2S04. These data indicate magnesium corrosion shows a negative difference effect also in acidic media where it is very unlikely for there to be any surface films. This indicates a reassessment is required of the current understanding of the mechanism of the negative difference effect. In addition, the magnesium dissolution rates and the hydrogen evolution rates in HCI were much faster than those in H$S04.

In contrast, Fig. 14 shows that there was no negative difference effect in 1 N NaOH, because no hydrogen evolution was detected in the anodic passivation region of magnesium over 24 h, whereas hydrogen evolution could be observed at a cathodic potential in this system. The magnesium dissolution rate in 1 N NaOH was extremely low, at all applied potentials in both the anodic or cathodic ranges.

. .) . Mg; HCI

- H,; HCI _ -* _ Mg; H,SO,

+ H,; H&30,

-100 0 100 200 300 400

Applied Current Density (mA/cmP) Fig. 13. Anodic magnesium dissolution rates and hydrogen evolution rates in 1 N HCl and in 1 N

HISO+

1992 G. Song et al.

-2000 -1500 -1000 -500 0

Applied Potential, E [mV] Fig. 14. Anodic magnesium dissolution rates and hydrogen evolution rates in 1 N NaOH

DISCUSSION

Reaction sequence./& magnesium corrosion

Our previous study I2 of the electrochemical corrosion of pure magnesium in 1 N NaCl found that the corrosion behaviour was consistent with a partially protective surface film. The surface film covered all the surface for applied potentials more negative than a critical potential E,, which was about 15 mV more negative than the corrosion potential E,,,,. Above Epl the surface film was only partially protective and its surface coverage decreased

with more positive applied potential, E. This prior study I2 provided strong evidence that a partially protective film plays a key role in the corrosion process of pure magnesium in NaCl solutions of pH 3-l 1. Furthermore, a partially protective film appears consistent with the

measurements in the present study for magnesium in Na2S04 solutions, in pH 13 solutions

and in NaOH solutions. The present study produced the data of Fig. 13 which showed that the hydrogen

evolution rate increased with increasing applied anodic current density also in 1 N HCl and 1 N H2S04. Thus, there was a negative difference effect even on a film-free magnesium surface. This indicates that the other three models mentioned in the introduction should be

considered for the magnesium corrosion mechanism when there is no surface film and also for the magnesium corrosion mechanism at imperfections or broken areas of the surface film on magnesium.

Of these three models, the univalent magnesium ion model is the most probable mechanism. It is able to explain the occurrence of the negative difference effect at a film-free surface as shown in Fig. 13 as will be discussed subsequently. The magnesium hydride model can not explain the increase of hydrogen evolution with increasing anodic potential or current density when there is no film on a magnesium surface.12 Furthermore, magnesium hydride is not stable and cannot exist in an acid solution, even though the overvoltage for hydrogen evolution on magnesium is as high as 1 volt.8*26 The particle undermining model


appears unlikely in acid solutions. The magnesium surface after corrosion in the acid solutions was smooth and bright with no indications consistent with particle undermining.

Therefore, let us assume that the univalent magnesium ion, Mg+, is involved in the anodic dissolution process at the broken area of a surface film (or at a film free area) on the magnesium electrode. There are a number of possible reaction pathways. The simplest possible reaction sequence is as follows.

K1+ Mg+E+Mg++e

I

Mg+ + H20 Kz\ Mg2+ + OH- + (1 /2)H2 (2)

2H20 + 2e % 20H- + Hz (3)

The first step in the anodic dissolution of metallic magnesium, reaction (1) produces Mg + ions by an electrochemical reaction. This step is assumed to be the rate determining step. K, + and K, _ are, respectively, the forward and backward reaction rates for reaction (1). Next, Mg+ is further oxidized to Mg2+ by the chemical hydrogen evolution reaction (2). K2 is the forward reaction rate for reaction (2). Since reaction (1) is assumed to be the rate determining step, (i.e. the slowest step), K I + , and KI _ should be much smaller than K2, and the concentration (C,) of the intermediate species Mg+ should be much lower than the Mg2+ concentration. It should be noted here that reactions (lH3) individually might be very complicated processes. Each of them may consist of many more detailed intermediate steps and include some other species which are not specified in (lH3) but can influence the rates of the reactions.

Let II and Zs be the current densities corresponding to the electrochemical reactions (1) and (3), respectively. Then:

II = FKI+ - FC,K,_ (4)

= Fkl+exp(E/pl+> - FGki-exP(-E/IL) (5)

13 = &exp(-E/P,) = 2FK3 = 2Fk3exp(-E/fli3) (6)

where, E is the electrode potential, F is the Faraday constant, kl+, kl _ and k3 are electrochemical reaction constants for reactions (1) and (3), respectively. PI + , PI _ and ps are Tafel constants corresponding to the reaction constants kl +, kl _ and k3.

The partial protective surface film still plays an important role in the magnesium corrosion process. The broken area of the surface film (or film free area) (0) increases with increasing applied potential or current density.‘T8-‘2 On the area (1 - 0) covered with a perfect surface film, there is cathodic hydrogen evolution (Zcp) and anodic magnesium dissolution (I,,). Therefore, the complete electrochemistry, the anodic and cathodic processes (Z, and ZJ of the magnesium electrode, is a combination of the reactions occurring at the broken area of the surface film and on the surface film and can be described by:

1994 G. Song et (11.

z, = ez,, + (1 - @I,, (7)

I, = ez,, + (1 - B)Ic, (8)

I,, and I,,, the anodic and cathodic electrochemical processes related to the film-free surface within the broken area of the surface film, can be expressed as:

LX, = II (9)

I,, = I! (10)

ICP and lap have a relatively small magnitude compared with II and I, because these reactions on a filmed surface are usually very difficult. Thus IcP and I,, can usually be neglected

whenever the surface film has any broken areas. I,, and Ia,, need to be taken into account only when the magnesium surface is completely covered by a perfect surface film.

For the case when there are some broken areas, the film free area can be expressed as 1 > 0 > 0. Since Ia, > > Zap, and I,, > > ICp, eqns (7) and (8) can be rewritten as:

Ia = 01, (11)

z, = ez, (12)

For the case of complete film coverage O-+0, eqns (7) and (8) become:

z, = zap

4 = Lp = ZFexp(-E/l&)

(13)

(14)

Thus the experimental data of Fig. 13 have led to a consideration of the reaction sequence on a film free magnesium surface or at broken area of surface film for magnesium corrosion.

The simplest possible reaction sequence is that described by reactions (lt(3). For this reaction sequence, current densities have been proposed for each individual reaction step and also for the measurable anodic and cathodic current densities. The next section compares the predictions from the proposed reaction sequence with the electrochemical impedance spectroscopy data.

Electrochemical impedance spectroscopy Electrochemical impedance spectroscopy (EIS) is a very useful technique for studying

electrochemical reaction mechanisms. An electrochemical corrosion system usually consists of several different. individual reaction processes. The different reaction processes have different reaction rates at a given electrode potential. Their responses to a changing applied electrode potential are also different. A reaction process produces a significant current only if the rate of the change of applied potential matches the rate of the process. Therefore, if a series of small AC potential changes with different frequencies are applied to an electrode, then the responses of the reaction processes involved in the electrode reactions occur in different frequency ranges. By analyzing these responses, the individual processes may be deduced.

The quantity measured in EIS is the electrochemical impedance of the electrode


(A) (a Ida X) (dX/dE) e 0 (B) (aI+ X) (dX/dE) > 0

-Zi HF,CP

Non Faradaic Faradaic

(dX,/dE) > 0 and

(dX,/dE) < 0

\LF,CP

Zr MFJD

LFJD

Fig. 15. A schematic theoretical Nyquist plots. HF = high frequency range; MF = middle frequency range; LF = low frequency range. CP = capacitive loop; ID = inductive loop.

interface, Z(w), which is the frequency dependent proportionality factor Z(o) = E(w)/i(o) between the excitation potential E= E, sin (wt) and the resultant current density i= i, sin (wt + $), where o is the excitation frequency and $ is the phase angle between the potential and current density. Z(w) is a complex-valued vector quantity Z(o) = Z’(w) + jZ”(o) with real, Z’(o), and imaginary Z”(w) components; and j = fl. A common data presentation is a Nyquist plot; -Z”(o) is plotted vs Z’(w) as shown in Fig. 15. The impedance would be equivalent to a simple resistance if the resultant current was in phase with the exciting potential; i.e. if $ = 0, then Z” = 0 and Z(w) = Z’(o) = R, where R is a pure resistance. This would be represented by a point on the Z’-axis. If the phase angle + = + 90”, then the impedance is equivalent to a capacitance with Z’ = 0 and - Z”>O. If however the phase angle 4 = - 90”, then the impedance is equivalent to an inductor with Z’ = 0 and -z”io.

The response of an electrode system, to a small AC perturbing potential signal, consists of two parts. One part of the response is a non-Faradaic current density AZNF which is a function of frequency. It is caused by the charging and discharging processes across the interface between electrode and solution. The other response is a Faradaic current density AZ,. It represents the change in the rates of the electrode reactions. Therefore, the interface is

1996 Cr. Song et (II.

equivalent to a parallel connection circuit consisting of a double layer capacitance Cdl and a Faradaic admittance YF. The admittance Y (or impedance 2) for the electrode is given by:

1 -=y=jwc& + YE z

(15)

Cao ‘3V’4 developed a theory for analysing the Faradaic processes for the case when diffusion was not involved in the reaction sequence.

If the electrode reaction is controlled by one surface state variable X, then Cao ‘3,‘4

showed that the Faradaic admittance YF can be expressed as

yr = l/R, + (aIFlax)(ax’/aE)jLjo - axfpx] (16)

where X’ = dXldt, and R, is the transient resistance defined by

R, = 8E/aIF (17)

Subsequently, Cao ” further developed the above expression to

YF = 1 /R, + (+/aX)(dX/dE)/[ 1 - jo/W’/aX] (18)

Correspondingly,

Y = j&d) + 1 /lit + (afF/aX)(dX/dE)/[l - jo/aX’/ax] (19)

Figure 15 (A) and (B) illustrates the predictions of eqn (19) on schematic Nyquist plots. There is in all cases as shown in Fig. 15 a capacitive loop in the high frequency range due to

the double layer capacitance, Cdl, as already discussed above. Figure 15 (A) shows that there is, in addition, a capacitive loop at low frequencies caused by the Faradaic electrochemical reaction process when (~~~/G’)(dX/d~~O. In contrast, Fig. 15 (B) shows that there is an inductive characteristic in the low frequency range when @Z~/aX)(dX/dE) > 0. ’

If there are two surface state variables that control the electrode processes, and‘if these two variables XI and X2 are independent of each other, then Cao ‘3,‘4.27 showed that the admittance of the electrode is given by:

YF =1/R, + (afF/aXl)(dXt/dE)/[l - jo/(aX;/ZJX,)]

+ (GPX2)(dX2IdE)I[t - jwl(aX;/aXz)l (20)

Combining the non-Faradaic and Faradaic processes leads to:

Y =jwCdl + I/ Rt + (alF/axl )(dXt /dE)/[ 1 - jo/(ax; lax,)]

+ G%/aX2)(dX2/d-m - j4(a$iax2)1 (21)

Figure 15 (C) and (D) shows schematic theoretical Nyquist plots for conditions corresponding to the experimental data of Figs 8 and 9; there are two capacitive loops and one inductive loop. The capacitive loop in the high frequency region is mainly caused by C,, and R,. In the middle frequency range (MF) and low frequency range, the two loops, one capacitive and the other inductive, are the contributions of the last two terms in eqn (21).

The electrochemical behaviour of the reaction sequence (l)-(3) can be assessed based on the analysis proposed by Cao ‘3,‘4,27 and summarised above. The reaction sequence (l)--(3)


involves two independent surface state variables: 0 which is the film free area of the surface and C,,, which is the Mg+ concentration. Let

xt = 0 (22)

x2 = c, (23)

The first variable 8 increased with increasing applied electrode potential E. That is, whenever 0 # 0, then

d0/dE > 0 (24)

The Faradaic current density is always given by the difference between the anodic and cathodic current densities, so that:

IF = Ia - I, (25)

When 8 # 0, according to eqns (11) and (12)

IF = B(Z, - 4) (26)

When magnesium was anodically polarised the measured current density increased steadily and monotonically with increasing potential, so that IF > 0, and Ii - Zs > 0, so the anodic current density should increase with 8. Thus

+/a0 = zl - z3 > 0 (27)

Combining eqn (27) with (24) predicts

@Zr/aQ(dO/dE) > 0 (28)

indicating the Nyquist plot should contain an inductive loop related to the surface state variable 0.

In contrast, for the cathodic polarisation branch, Zr < 0, or Ii - Zs < 0, hence

az,/ae = zI - I3 K 0 (29)

(aZr/ae)(dO/dE) < 0 (30)

and a capacitive loop caused by 8 variation would be expected. In summary, eqns (28) and (30) predict that the first surface state variable 6, the broken

area of the surface film, produces an inductive loop in the EIS diagram for anodically polarised magnesium, and gives a capacitive loop for cathodically polarised magnesium. These predictions are for a partially protective surface film. If the surface film is perfect, that is if 8 = 0, then there will not be these loops.

For the second variable C,, in the case of 8#0, substitution of eqn (6) into eqn (26), leads to:

IF = 8(F&+ - &f-f& - 13) (31)

which implied that:

1998 G. Song et al.

azFlac, = -0FK,_ < 0 (32)

Since the AC impedance experiment was conducted when the electrode system was at its

steady state, then

dC,ldt = 0(K,+ - C’,K, - C’,,,Kz) = 0 (33)

SO

dC,ldE = [(K,,IP,+)(KI + Kz) + KI,~-IILII(KI- + Kd2 (34)

Simply supposing the transferring parameters for the reactions (l)-(3) were the same, at

about 0.5, then f12 = j3, = PI + = p, then eqn (34) can be rewritten as:

dC,ldE = [(KI+(KI~ + K2) + Kid-I/[WL + &)I’

= WI+KI + KI.+KMP(KL + KI)]’ (35)

Therefore

dC,/dE > 0 (36)

so.

(a/,/aC,)(dC,/dE) < 0 (37)

Equation (37) means that the intermediate species Mg ’ produces a capacitive loop in the EIS diagram if the magnesium surface film is not perfect, that is if 8 #O. However in the case of a perfect film, i.e. for 8 = 0, eqn (31) is no longer valid. No C, should be considered and IF

has nothing to do with the C,,. Thus there will be no capacitive loop caused by C, in the EIS

diagram. The above analysis of AC impedance based on the Cao’s theory ‘1,‘4,27 can successfully

explain the Nyquist plots of Figs 8 and 9. The anodically polarised magnesium electrodes in both NaCl (Fig. 8) and Na2S04 (Fig. 9) exhibited one capacitive loop and one inductive loop in the middle and low frequency ranges. The inductive loop is attributed mainly to the

partially protective surface film and the capacitive loop is related to the MgC ion concentration within the broken area. In the NaCl solution, the presence of Cl- would

cause the surface film to be much more active than in the Na2S04 solution, so that those electrode processes in NaCl related to the surface film would be much faster than in Na,S04. As a consequence, the time constant of the AC response process related to the surface film was much smaller in NaCl, and the corresponding inductive loop appeared in a higher frequency range in the Nyquist plot. That is why in NaCl, the inductive loop was in a middle frequency range and the capacitive loop in the lower frequency range, while vice-versa in Na2S04 (see Figs 8 and 9).

When magnesium was cathodically polarised, there were no capacitive and no inductive loops in the intermediate and low frequency range, in both NaCl and Na2S04 (Figs 8 and 9).


This is consistent with complete film coverage of the magnesium surface at a sufficiently negative cathodic potential.12 In this case, 6 and C, were no longer surface state variables. The perfect film acted only as a parallel connection circuit across the interface equivalent to a resistance (R,) parallel with a capacitance (Cdl). Therefore there was only a capacitive loop in the Nyquist plots (Figs 8 and 9).

The Nyquist plots for magnesium at J&r, p resented a case intermediate to those at anodic and cathodic potentials. At E,,,,, 8 and C, were still surface state variables, so the Nyquist plots were similar to those of anodically polarized magnesium in both NaCl and Na2S04 (Figs 8 and 9).

In the high frequency range, there was always a capacitive loop in the Nyquist plots (Figs 8 and 9). It was caused by the Rt//Cdl equivalent circuit. The increase of Cd, with increasing potential (see Figs 10 and 1 1), was caused by the increase of the broken area of the surface film with increasing applied potential, and because the broken area had a much larger capacitance than the film covered area. The decrease of Rt with increasing applied potential in the anodic region (Figs 10 and 1 l), has a similar explanation. The broken area of the surface film increased with increasing potential and the reaction resistance of the broken area was much lower than that of the film-covered area. Rt had a maximum value at a cathodic potential close to the corrosion potential, and Rt deceased for more negative cathodic potentials. There was an almost perfect film when the potential was sufficiently negativeI and the only reaction on the surface was hydrogen evolution (I&. According to (15), the reaction resistance for hydrogen evolution on the film should be RtH = l/(al,/aE) = &..Zc. Thus Rt should decrease with decreasing potential because 1, increased with decreasing potential.

The above discussion indicates that the proposed model (lHl4) involving Mg+ is in agreement with the experimental results.

Negative difference effect If it is assumed that the reactions on magnesium through a perfect film can be neglected,

and if the surface film is only partially protective, then the total magnesium dissolution rate, V,,,, is equal to the rate of reaction (1):

V, = (VI+ - G-C,) = Wi+exp(EIP,+) - C&-ev(-E/P,-)1

Similarly, the total hydrogen evolution rate, Vi.,, is the sum of reactions (2) and (3)

vi, = 9C,,,(K2 + K3)

which, incorporating eqns (6) and (33), can be rewritten as

vh = e(K2 + K3)Kl+/(Kl- + K2)

= %+exp(E/P,+)W2 + k3exp(-EIP3)ll[kl-exp(-E/P1_) + K21

(38)

(39)

(40)

Equations (38) and (40), show that the broken area 8 of the surface film plays a very important role in determining the magnesium dissolution and hydrogen evolution rates. In other words, when there is a partially protective surface film on magnesium, the coverage of the film is mainly responsible for the anodic dissolution behaviour and the negative difference effect, because the increase of 8 with E causes the simultaneous increases of V, and vh. This is the idea that was employed to successfully explain the experimental results

2000 G. Song el ul.

for magnesium dissolution and hydrogen evolution in NaCl solutions in our previous

work.12 Similarly, 8 should be the main factor controlling the anodic magnesium dissolution and hydrogen evolution rates at different applied current densities measured in the present

study and shown in Fig. 12. The data of Fig. 12 show that the magnesium dissolution rate was reduced to almost zero for an applied cathodic current density of - 1 mA/cm2 in both NaCl and Na2S04 solutions. This implies nearly perfect film coverage, i.e. 8+0 at that applied cathodic current density.

Further strong evidence supporting the partially protective film model was shown by the behaviour in 1 N NaOH. Under these conditions, Mg(OH)2 was very stable and as a consequence the polarisation curve of Fig. 5 showed that magnesium was clearly in a passive state. This implied nearly perfect film coverage, i.e. e-+0. Therefore there was the

expectation of almost no magnesium dissolution and no negative difference effect, consistent with the actual experimental observations of Fig. 14.

In contrast, when 8-t 1, for example in H$04 or HCI acid solutions, the magnesium dissolution and hydrogen evolution rates can be simplified to:

C’, = KI, - KI Cm = kl+exp(EIP,,) - C&-exp(-El(3,-) (41)

vh = kl+exP(E/h+)[K2 + k3exp(-EIB3)ll[k~~exp(-EISI) + &I (42)

Equation (41) predicts that the magnesium dissolution rate (I’,,,) increases with increasing applied potential if the applied potential is higher than the equilibrium potential (Elo) for

reaction (1). This is completely understandable as the standard equilibrium potential for reaction (1) is more negative than -2600 mV/scn,8.‘2 and is consistent with the experimental results as shown in Fig. 13. In fact, during our experiment, all the applied potentials were much higher than Et”, and k, exp (-E/p, -.) was always very small compared with the other terms and could be neglected in expressions (38)-(42) so we also

have:

VITl = kl+exp(ElIJ ,.) Q-+1 (43)

vh x (kl +lKdexp(EIh +)k + ksM-EIP3) 1 e-9 1 (44)

v ,,, ̂ ” 0kt+exp(EIPI+) 1>e>o (45)

Vh x e(kI+lK2)exp(EIP,+)[K2 + k3ev-EIPd 1>e>o (46)

Now let us discuss the hydrogen evolution. According to eqn (42) or (46), even in acidic solutions when 8 + 1. vh should also increase with increasing potential. This prediction explains the negative difference effect of Fig. 13.

The rates of magnesium dissolution and hydrogen evolution in the anodic range in Cl ~~ containing solutions were higher than in the Cl ~ free solutions. Explanations can be related to different values of surface film coverage, 8, and to different rates of reactions (l)-(3) in the different media. It is reasonable to believe that the broken area of the surface film was larger in NaCl than in Na2S04. Then, eqns (38) and (40), or (45) and (46) predict that the magnesium dissolution and hydrogen evolution rates in NaCl should be higher than in Na2S04, consistent with the experimental results of Fig. 12. However, there is also the


possibility that Cl- changed the rates of reactions (l)-(3). Evidence for this comes from the experiments of Fig. 13 in HCl and H$SOa when there was no surface film. The magnesium dissolution and hydrogen evolution rates in HCl were much higher than those in H2S04. This suggested that Cl- did affect the rates of reactions (l)-(3). In particular, the magnesium dissolution rate in HCl which was nearly double that of in H2S04 (Fig. 13) implied that, according to equation (41) or (43), k, + was dramatically higher in HCl than in H2S04. If k, + is increased, then Vi, should also be increased according to eqn (42) or (44). This is consistent with the observation that the hydrogen evolution in HCl was much higher than in HzS04 (see Fig. 13).

The influence of Cl- on ki + and the activity of the surface film can arise from the competitive adsorption of Cl-, OH- and S042- on the magnesium surface. It is expected that the Cl- is much more aggressive, whereas the other two, especially OH-, might have some inhibiting effect when they are adsorbed on the electrode. In the NaCl solutions, a great amount of Cl- would be expected to be absorbed on the surface film and on the film- free surface. Perhaps the Cl- ions not only made the surface film, magnesium oxide or magnesium hydroxide, much more soluble, but also combined with metallic magnesium, weakening the bonds among the metallic atoms and accelerating their departure from their original position to form the intermediate species Mg + .

Polarisation behaviour Based on the above model, the current density of polarisation curves for a magnesium

electrode with a partially protective film, Z= Z, - Z, = 8(Zi -I& is given by

Z = eF[kl+exp(EIBi+) - G&-exp(-E/Pi-) - 2k3exp(-EIB3)1 (47)

After the term of ki _ exp( - E/S, _) is neglected, this can be simplified to

Z = Wkl+exp(EIPi+) - 2k3ew(-ElP3)1 1re>o (48)

For the case of a perfect surface film on magnesium, the polarization curve can be formulated as

Z = Zap - Z,exp(-E/P& O-+0 (49)

For magnesium in NaCl and Na2S04 solutions, some broken areas of significant size were observed at applied potentials E > Ept which was always slightly more negative than EC,,, (see Fig. 2). The decrease of Ept with increasing Cl- content shown in Fig. 2 implies that Cl- decreased the stability of the surface film and initiated film breakdown at a more negative potential. An increase in SOd2- content had an effect reverse to that of Cl-. The reason might again be the competitive adsorption of these two ions.

In the cathodic potential range where E< Ept, the surface film was regarded as nearly perfect, 8 + 0.12 In this cathodic potential range, Z,, could be neglected, and the cathodic current Z from equation (49) can be rewritten as

Z = -Z,exp(--E/P&, E c Ert, (50)

The above equation predicts Tafel behaviour for the cathodic branch of the polarisation curve for E < Ept consistent with the measurements of Fig. 1. Figure 3 shows that b, (b,, = 2.3@,) does not change much with changes of Cl- and SOd2- content, indicating that the

2002 G. Song er (11.

mechanism of cathodic hydrogen evolution was almost independent of Cl- or SOe2

content. However, Fig. 4 shows that I,,, the exchange current density of hydrogen evolution on the surface film decreased with increasing SOd2- concentration. The measured trends (Fig. 1) in the dependence of cathodic current density on solution composition were

consistent with the measured trends for h,, and for I,,. The possible reasons for the changes of I,,, might be that the activity of the surface film was increased by decreasing SOd2- content or increasing Cl --~ content, or the thickness of the film decreased with increasing Cl ~- concentration or decreasing SO4’ _ concentration.

The stability of a surface film also depends on the solution pH value. A higher pH value solution usually makes the surface film more passive. In this solution, there is a greater

concentration of OH- available to compete with Cl- ions to absorb onto the film surface, and to repel more Cl -- ions out of the film/solution interface. This decreases the influence of

the Cl- ions on the film. This is consistent with a smaller influence of Cl- ions on the

cathodic branch of the polarisation curve in the higher pH media (curves C and D at pH 13. Fig. 1) compared with the influence in a lower pH solution (curves A and B at pH 11, Fig. l),

and almost no significant influence in 1 N NaOH (Fig. 5). When E > Ept, O>O, and the current density is given by eqn (47) or (48). Equations (47)

and (48) indicate that the anodic branch of the polarisation curve is much more complicated

than the cathodic branch because of the second variable 0. Furthermore, the anodic branch

of the polarisation curve in Cl ~ containing solution should be much higher than in the Cl-

free solution (Fig. I), because increasing Cl ion concentration increases both the magnitude of the constant ki + and the broken area (i.e. increased 0).

The surface film also plays different roles in the anodic and cathodic processes. At an

applied cathodic potential within the Tafel region where 0 -+ 0, hydrogen evolution occurred on the film surface. The presence of Cl ~ ions might only make the film thinner with hydrogen evolution still occurring on the film surface. In contrast, the magnesium dissolution into solution at an applied anodic potential occurred mainly at the broken area of the film. Increasing Cl ~ caused an increase in 8 and provided a larger area for anodic dissolution. Therefore. the anodic process rate was much more sensitive to Cll concentration than the cathodic process, so that the increase of the anodic branch of the polarisation curve was higher than that of the cathodic curve (see Figs 1 and 5). If there was no surface film on the magnesium, then the above arguments would not be valid as, e.g. on HCl and HzS04 (see Fig. 7).

In fact, the anodic magnesium dissolution became significant, and the anodic current density could be measured only after the surface film became partially protective. Therefore the breakdown of the surface film immediately caused the change of current density from cathodic to anodic, and Ept occurred only slightly earlier than E,,,, (Fig. 1). This explains the close coupling of En, and Ecorr, (that is Ept was only about 15 mV more negative than E,,,,) and explains why both quantities showed similar changes to changing Cl- and S042- concentrations (Fig. 2).

Figure 6 shows some special cases. In 1 N NaOH no pitting was observed. The surface film can be regarded as perfect irrespective of the Na2S04 concentration. In this case, eqn (49) can be rearranged to give the following expression for E,,,,:

E corr - Pp,W,lL,) (51)

In contrast to the situation in solutions of lower pH (11 or 13), the surface film was very


stable in the much more basic solution like 1 N NaOH. In such a solution SOd2- should have little influence on SPc, Ipc and lap, or the addition of SOd2- might make the surface film more stable and decrease Zap, so there was only a small increase in Ecorr. For small concentrations which is not high enough to cause pitting corrosion in surface film, the influence of Cl- on E,,,, was small. In contrast, when Cl- content was increased to a high concentration, pitting appeared on the surface at E,,,,, and it was found that E,,,, decreased with increasing Cl- concentration (see Fig. 6).

In the vicinity of E c,,rr, eqn (48) can be rearranged to provide the following expression for the polarisation resistance R, for magnesium covered with a partially protective surface film:

l/R, = l/R,- = (dl/dE)IE+,, = Wrl+exp(K%+) - 2k3exp(-E/P~ll(de/dE) + el~~+ew(E/B~+M%+ + 2k3exp(-E/P3)/P31)E=E,,

= F(kl+exp(E/P,+)[(de/dE) + (e/h+)1 (52)

+ 2hw(-E/hNe/iQ - (de/W1h-,,9 be>0

Equation (52) predicts that R,_ decreased with the increasing Cl- concentration, because increasing Cl- concentration increases 8 and kl +; this prediction is consistent with the experimental measurements shown in Fig. 2.

CONCLUSION

A partially protective surface film plays an important role in the electrochemical dissolution processes for magnesium in NaCl, Na2S04 and NaOH solutions.

In the broken areas or on a film-free surface, the experimental data are consistent with the involvement of the intermediate species Mg + in the magnesium dissolution process. Magnesium is first oxidised to the intermediate species Mg + , then the intermediate species chemically reacts with water to produce hydrogen and Mg2+ .

The presence of Cl- made the surface films more active or increased the broken area of the film, and also accelerated the electrochemical reaction rate from magnesium to magnesium univalent ions.

Acknowledgemenrs-The authors wish to thank Prof. H.C. Lin and John Oweczkin for their support in the EIS experiment and ICPAES analysis.

REFERENCES

1. R. Tunold, H. Holtan and M.-B. Hagg Berge, Corros. Sci. 17, 353 (1977). 2. E. Gulbrandsen, Electrochim. Acta 37(S), 1403 (1992). 3. W.S. Loose, Corrosion and Protection of Magnesium, ed. L.M. Pidgeon, J.C. Mathes, N.E. Woldmen et al.

ASM, 1946, pp. 173-260. 4. I.J. Polmear, Light Alloys, Metallurgy of the Light Metals, 2nd Edn. Edward Arnold, 1989. 5. L.L. Shreir, Corrosion-Metal/Environment Reactions, Vol. 1. Newnes Butterworths, 1965, pp. 86100. 6. E.F. Enket, Principles of Magnesium Technology, Chap. IX. Pergamon Press, 1966. 7. N. Pebere, C. Riera and F. Dabosi, Electochim. Acra 35(2), 555 (1990). 8. G.G. Perrault, in Encyclopedia of Electrochemistry of the Elements, Vol. VIII, ed. A.J. Bard. Marcel Dekker,

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