Galvanic Corrosion-A
Kinetic Study
T. H. Randle Swinburne University of Technology, P.O. Box 218, Hawthorn, Victoria 3122, Australia
Electrochemistry is a common component of both beginning and advanced chemistry courses. The corrosion of metals in aoueous solution is a n electrochemical ohenomenon of substantial economic significance. Galvanic corrosion-the enhanced corrosion resulting from the direct contact of two or more dissimilar metals in an aqueous env i r o n m e n t i s a common t v ~ of e corrosion. This exoeriment provides a comprehen&e investigation of the !&etics of galvanic corrosion.
Galvanic corrosion occurs when two dissimilar metals are electrically coupled, for example, placed in direct contact. In Figure 2 the dashed lines represent the currentpotential relationships for the anodic and cathodic reactions occurring on each metal in the isolated (uncoupled) condition. The full lines show the total anodic and cathodic currents a t any potential for the coupled metal system.
Theory The four primary steps involved in metal corrosion are shown in Figure la. (1)Metal oxidation (dissolution)at an anodic site (2) Reduction of an oxidizer (Ox)in solution at a cathodic site
on the metal surface (3) Ian migration within the solution to maintain electrical (4)
neutrality throughout the solution Electron transfer within the metal between the anodic and cathodic sites
The four steps are series-linked. Thus, the rate of metal corrosion (step 1) will be the same as that of the slowest step, usually one of the two electrochemical steps (1or 2). Because each step involves movement of charge, cormsion rate may be defined as a current I,,. Because two charpe-transfer reactions occur simultaneously on the metai surface, the uniform potential adopted b; the corroding metal Vwm,a s measured against a designated reference electrode. will be a mixed ootential. Its value will be somewhere between the equilibrium potentials for the anodic and cathodic reactions. In corrosion literature it is common to represent the current-potential characteristics for the two electrochemical reactions on one set of axes to give an Evans diagram (Fig. lb). The more complex situation of galvanic corrosion is illustrated in Figure 2. The following assumptions are made. Metal 2 is less active toward corrosion than metal 1
ve,,
< ve,
The same cathodic reaction occurs on both metals.
Ox+ne+R T h e area of metal 2 is greater than that of metal 1, and the kinetics of the cathodic reaction are more favorable on metal 2 (larger cathodic currents at any potential)
I
v 'corr
Current (I I
Figure I.(a) Simplified metal corrosion in aqueous solution. (b) Polarization curves for the oxidation and reduction reactions occurring on a corroding metal: an Evans diagram.
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a Recorder
& reduction r a t e /
Figure 2. An Evans diagram tor galvanically stimulated corrosion resultino from the couolina of two dissimilar metals in which metal 1 is moreactwe than mdta 5 Dashw I nes represent po arlzauon cLrves for reactions occdrr ng on eacn metal in the unwbpled slate FLI mes represent Iota anodlc an0 catnodlc current at any potent al for the coupled metal system. Thus, for the coupled system the total corrosion current is given by I& a t the corrosion potential V&. The cornsion rate of each metal is now given by the rate of the anodic dissolution reactioa M+Mnt+ne
in Figure 2. As a t the potential V& represented by indicated, the more active metal now corrodes a t a greater rate than in the uncoupled condition, The difference between the two corrosion currents is the galvanic corrosion current I,. Conversely, although not indicated directly in Figure 2, the corrosion rate of the less active metal is decreased in the galvanic situation relative to that in the uncoupled condi-
This of course is the principle of cathodic protection. It is the more active (i.e., more reducing) metal that satisfies the increased demand for electrons due to the higher rate of the cathodic reaction when the metals are coupled. I t should be noted that Igis not the total corrosion rate of It only the active metal in the coupled condition appro:?mntes the total rat@when the area of the more stable metal is much greater than that of the active metal. Experimental Electrodes and Cells Figure 3 shows the experimental setup in block form. The galvanic couple chosen was zinc and copper. Planar electrodes of each metal were cut from foil. After abrasion and cleaning, the single exposed surface of each electrode was defined with stopping-off lacquer. The active area of
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.Figure 3. A block diagram of the zero-resistanceammeter and galvanic cell for the zinclcopper couple. the zinc electrode was 12.5 cm2, whereas the various copper electrodes used throughout ranged in area from 1cm2 to 50 cm2. The electrolvte was 3.5% sodium chloride solution (artificial seawater; pH = 6.4). The cell was a simple beaker (800 cm3)thermostatted at 25 T in a water bath. Eledrolyte stirring, when used, was by motor-driven glass propeller. Measuring Current Galvanic current cannot be measured accurately by a simple ammeter because the finite resistance of the arnmeter a d s to decrease the current being measured. This disadvantage is overcome bv the zero-resistance ammeter (ZRA). ~ k a n u a l l ybalanced ZRAmay be constructed from simple circuit elements ( I ) . A convenient, auto-balancing ZRA is easily constructed from a n operational amplifier (2) in which the am~lifieris used in the current-follower configuration (Fig. 3). The OD am0 (741f 15V)keens the notential difference a t zero beiweek its input terminks, acverified continuously by the high impedance (>lo8R) voltmeter; current is passed through the feedback loop that contains the cell and the current-measuring resistor (1-100 R). Because the net current a t the summ&g point S must be zero (Kirchoff's law), we get
Hence measurement of V, [output potential of op amp, allows determinati~nofl,~, becauseR is set. 1,ll is the current when the potential dfferencc between the electrodes is zero [the equivalent of jialvonic coupling,, so I,,ll = I,. A potentiometnc recorder is ideal for continuous monitoring of I,. Other equipment used included a pH meter, an oxygen electrode (Clark type, voltammetric detector), and a conductance bridee. All emenmental electrode notentials N . ). were measured with respect to a saturated calomel electrode (SCE) (3) and are quoted as such.
-
Mass Transport
Such a relationship is typical of a process controlled by the rate of diffusion (mass transport) of a reactive solution species to the planar electrode surface (4). The rate (current) tends to a steady state value rather than zero for long periods (e.g., greater than 20 s): Natural convection due to density gradients in the solution pegs the thickness of the diffusion layer (depletion zone) a t a constant value (5). Stirring the Electrolyte The presence of mass transport control is further verified by stirring the electrolyte. I, will immediately increase above the value in unstirred solution because the rate of mass transport of the active species is increased by the forced convection. If solution stirring is smooth (laminar solution flow), then I, increases approximately linearly with the square root of stirring rate. (Linear variations were obsewed in this laboratory up to 2,500 rpm). By contrast, the rate of an electrochemical reaction proceeding under charge transfer control (i.e., electron transfer across the electrode-solution interface) is independent of time and electrolyte agitation. Concentrations ofReactive Species As well as being sensitive to stirring, mass transport controlled currents vary linearly with the concentration of the reactive species in solution. For metallic corrosion in chemically simple, aqueous electrolytes the active oxidizer is usually O2 or Hi(aq) (or both).
As shown in Figure 5 a linear relationship is observed between oxygen concentration in solution and I,, indicat-
Fioure 4. he variation of oalvanic current densitv for the ZnICu mupli; in unstrred 3 5% ~ a ~ - s o i ~ (a) t o wn rh t#m&andibj with the nverse ot the square root of time. Results
Zn is the more active (anodic) metal in 3.5% NaCl solution, which is easily verified by measuring the corrosion potential of each metal. Typical values are
V,,, for Zn = -1.053 V V,,
for Cu = -0.200 V
(This result is predicted fromE" values although this is not always possible. For example, compare Zn and A1 in 3.5% NaCl solution.) After identifying Zn as anodic and Cu as cathodic in the galvanic couple, important questions remain. Which of the four steps in Figure l a is the rate-determining step for galvanic corrosion? What factors influence the rate of corrosion? The Rate-Determining Step
The obsewed variation of I , with time (t) in unstirred electrolyte is shown in Figure 4a. The actual relationship, a t least fort < 20 s, is
where k is a constant (see Fig. 4b)
Relative oxygen concentration Figure 5. The dependence of galvanic current density forthe ZnICu couple on relative oxygen concentration in a stirred 3.5% NaCl solution (rotation rate of stirrer: 1,000 rpm). Volume 71 Number 3 March 1994
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ing that reduction of Oz a t the copper cathode is the ratedetermining step for the galvanic corrosion of the ZdCu couple in 3.5%NaCl solution. In Figure 5 the natural concentration of O2 in solution a t 25 'C is given the relative value of 1.(The actual concentration of Oz in water a t 25 'C is 2.6 x lo4 M (6)). Relative oxygen concentrations greater than 1 were achieved by passing oxygen gas through the solution. Both the oxygen content (oxygen electrode) and the corresponding galvanic current were determined a t various stages during oxygen sparging. Relative oxygen concentrations less than 1were achieved by passing nitrogen gas through the solution. aeain notinel. and oxygen - - content a t various times during nitrogen In Fimre 5 the deviation from linearity a t low Oz wncentrati0n.s may indicate a small contributibn of HYaq) reduction to the cathodic current. The significanceof oxygen can also be quickly demonstrated by a&ng oxygen scavengers (e.g., NazSOsand Nz&) to solution, causing an immediate decrease in I, This is a good place to introduce the topic of corrosion inh~bitorsif desired. Variations in pH Although the thermodynamics favors oxygen as the domthe rate of Ht(aq) reduction, inant species (greater P), being pH-dependent, is greater than that of Oz reduction a t least in acid solution (7).Thus, it might be anticipated that the rate-determining cathodic step will change with solution pH. The variation in I, with solution pH was examined over the pH range 12 to 1by stepwise addition of NaOH or HC1 to the cell electrolyte. Due to passivation, the influence of pH on I, is complex a t high pH (> 9)because it is dependent on time and the direction of pH change. (This influence will not be discussed further here.) However. with a decrease in DH(7+ 1). the chanee in I. is quite marked and uncompl~cated.~ l t h o u g hI, remain: reasonablv constant between DH 7 and 3. it increased sharply b i o w pH 3, rising by a Factor of about 50 between DH3 and 1. The increase in I, is due to the increasing contribution of H+(aa)reduction to the cathodic current. Hz bubbles becomeapparent on the electrode surface a t lower pH. In ad-
-
1
V)
I
C
I
C a,
0.4
a
1.2
4-
2 1.8 C
L
3
u
1.2D i s t i l l e d water
Fig~re7. Tne variat'on of galvanic cLrrent density with ~nlerelenrode separat on fortne Zn CL WJP e IF tap water and d st1 led waler (area of Zn and Cu electrode: 12.5 cm') dition, due to the relatively high concentration of Ht(aq) in solution at pH 1,the galvanic current is no longer stirringdependent because the galvanic corrosion process is no longer controlled by mass transport. Factors lnfluenclng the Rate of Galvanic Corrosion Apart from oxygen concentration, pH, and agitation of the electrolyte, other factors influencing the rate of galvanic corrosion are cathodic area and electrolyte resistance. Cathodic Area
20
40
60
Cathode area (cm2) Fiaure 6. The effect of cathode area on the aalvanic current densitv ofihe ZnICu couple in stirred 3.5% NaCl &lution (rotation rate df stirrer: 1.000 rpm).
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Journal of Chemical Education
Electrochemical reactions are heterogeneous, their rate being directly proportional to the area of the surface on which they occur. Figure 6 illustrates that the expected linear dependence of I, on cathode area is indeed observed, even when the cathode area is 4 times that of the anode. This auantitativelv illustrates the oft-auoted adaee that when the joining i f dissimilar metals is unavoidakle the area of the more active metal should be much ereater than that of the less active metal to avoid catast~ophiccorrosion. Electrolyte Resistance Ionic migration is an essential process in a corrosion cell (Fig. lb). Thus, it might be anticipated that electrolyte resistance will influence I, Electrolyte resistance may be
varied by altering the interelectrode distance (I)or the electrolyte (e.g., to tap water or distilled water). The observed variation in Zg with electrode separation in tap water and distilled water is given in Figure 7. No dependence ofIgo n 1 (1- 9 cm) was observed in 3.5% NaCl solution, presumably because its specific resistance (p) is too low (21R cm) for ionic migration to be kinetically limiting. In addition to the much smaller galvanic current observed in tap and distilled water relative to that in 3.5% NaCl solution, the current was also independent of electrolyte agitation and time, indicating absence of mass transport control. In distilled water (p = 1.0 x 10' R cm) the linear dependence ofIgon the inverse of electrode separation is rationalized as follows. For an ohmically conducting solution,
and
In this work AVd = 0.823 V in distilled water. The nonlinear variation of I, with I in tap water indicates that the specific resistance is too low (p = 1.6 x lo4 ~2cm) for ~ O N C migration to be completely rate-determining. Conclusion Through the subject of galvanic corrosion this experiment may be used to expose students to theoretical, experimental, and instrumental aspects of electrochemistry in an integrated, problem-solving context. The experiment has been carried out successfully for several years by senior undergraduate students and participants in a short course on corrosion. Literature Cited 1. Baboian. R. In Elecfmcherniml Techniouss for Cmsion: Bahian. R.. Ed.: NACE:
...
where AV.,, is the potential drop through the solution between the two electrodes;A is the electrode area; and R,,I is the solution resistance. Thus, if AV,.l is constant with electrode separation, then
tions: Wile": New Yark. 1980: oo 1 3 6 1 4 . 6. ~ o c k r k : 0%; ~ . Reddy A. K N. Modem Eleetmehemisfry;Plenum:New Ymk, 1973; pp 1050-1052. 6. CRC Handbmk of Chemise nnd Physics, 6ah ed.; Weast, R. C., Ed.; CRC Press: Boes Ratan, FL, 1979: p B-104. New York,1973; 7. Boek&,J. O'M; Reddx A. K N. ModomEleetmhem~f~;PI~~~m: p 1298. ~~
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