Ind. Eng. Chem. Prod. Res. Dev., Vol. 17,No. 2, 1978
(Silverwood e t al., 1962) and the substituted thiophenes (Rasmussen e t al., 1946). The types of sulfur products will be the subject of another investigation.
Acknowledgment The authors wish to record their appreciation to the Kuwait National Petroleum Company (KNPC) and the Kuwait Oil Company (KOC) for providing the asphalts and sulfur. Literature Cited Corbett, L. W., Anal. Chem., 36, 1967 (1964). Corbett, L. W., Anal. Chem., 41, 576 (1969). Grant, F. R., Hoiberg. A. J., Proc. Assoc. Asphalt Paving Technol., 12, 87 (1940). Haley, A. G., Anal. Chem., 47, 2432 (1975)
169
Hughes, E. C.. Hardman, H., Proc. Assoc. Asphalt Paving Technol., 20, 1 (1951). Jewell, D. M., Albough, E. W., Davis, B. E., Ruberto, R. G., lnd. Eng. Chem. Fundam., 13, 278 (1974). Marcusson, J., Z.Angew. Chem., 29, 21 (1916). Petrossi. U., Bocca, P.L., Pacor, P., lnd. Eng. Chem. Prod. Res. Dev., 11, 214 (1972). Rasmussen, H. E., Hansford, R. C., Sachanen, A . N., lnd. Eng. Chem., 38,376 (1946). Rostler, F. S.,Sternberg, H. W., lnd. Eng. Chem., 41, 598 (1949). Silverwood, H. A,, Orchin, M., J. Org. Chem., 27, 3401 (1962). Traxler, R. N., Schweyer, H. E., OllGasJ., 52, (19), 158 (1953). Tsurgi, J., Rubber Chern. Technol., 31, 762 (1958). Tucker, J. R., Schweyer. H. E., Ind. Eng. Chem. Prod. Res. Dev.. 4, 51 (1965).
Received f o r review September 6, 1977 Accepted January 23,1978
Symposium on Interfacial Phenomena in Corrosion Protection A Review of Electrochemical Corrosion Fundamentals Thaddeus M. Muzyczko The Richardson Company, Research and Development Division, Melrose Park, lllinois 60 160
Of the many types of corrosion, those brought about by electrochemical reactions are a cause for major concern and study. For this type of corrosion to occur a corrosion cell consisting of an electrolyte, a cathode, an anode, and a flow of electrons between these electrodes is needed. The basic thermodynamics and kinetics are reviewed using “ideal models”. Limitations of these models and practical extensions for real cases are discussed. A summary of what we do know and do not know about electrochemical corrosion is presented.
Introduction The corrosion of metals is a multibillion dollar worldwide problem that results in losses of materials, energy and even lives (Znd. Week, 1975). If you own a car and drive in a metropolitan area, corrosion is a readily visible, annoying phenomenon. Pitted chrome trim, rotting rocker panels, and frozen bolts are loud testaments to the fact that we have not yet licked the corrosion problem. T o be perfectly fair, the solution to corrosion problems requires many considerations these days: effectiveness, ecology, environment, effluent, emissions, energy, economy, and exasperation (Mazia, 1977). Often post-treatments and/or periodic retreatments of surfaces are the only current answers. Corrosion in general may be classified by its appearance, uniform or nonuniform, which may be microscopic or macroscopic (Henthorne, 1971a).The nature of the corrodant (wet or dry) is also a means of classification (see Figure 1). The two major mechanisms of corrosion are by direct chemical and electrochemical reactions. Most corrosion reactions, particularly of iron, are electrochemical in nature. Electrochemical corrosion is here defined as the unwanted and usually destructive oxidation of metals. I t is the main topic of this paper. Major considerations are surfaces and interfacial phenomena. For electrochemical corrosion to occur, a corro0019-7890/78/1217-0169$0.100/0
sion whole cell is needed, Le., an electrolyte, electrodes (anode-cathode), and an electron flow between the anode and cathode. Attempts to stifle these factors are the basis of current attacks on corrosion.
Ideal Models The concept of a so-called “ideal model” will be used as a convenient basis to introduce the well-studied corrosion reactions. The more complicated and unfortunately more common “nonideal models” will be introduced later. An ideal model is a well-behaved, uniformly, freely corroding metal in contact with an ideal electrolyte and having no inhibiting or accelerating side reactions. A representation of such a surface is shown in Figure 2. For corrosion to occur as with any chemical reaction, two major questions arise: (1)will a metal corrode in a given environment to a calculable equilibrium point? and (2) how fast? The first question is answered by thermodynamics; the second by kinetics. Thermodynamics Why do some metals corrode in given environments? Can their tendencies to corrode be predicted? Thermodynamics can answer these questions. Practical observations indicate that most commercial metals are more stable as oxides or compounds rather than in their free metallic forms. Iron is 0 1978 American Chemical Society
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$\
Examples: Fe I S O 1 Fe * 02
ti? T
Table I. Free Energies of Formation AG"f for Some Metal Oxide Corrosion Products AGOf, kcal
Compound
AGof, kcal
mol-'
Compound
mol-'
ZnO PbOz NiO CUO AgaO Au20sa
-76.1 -52.3 -51.7 -31.4
A1zOda) -376.8 Crz03 -259.2 FesO4 -242.4 Fen03 -177.1 MgO -136.1 SnOz -124.2 Does not form spontaneously.
-2.6
+29.1
Figure 1. Classification of metallic corrosion. where T = 25 OC, P = 1atm, M = a metal, c = crystalline solid, and g = gas. The following Gibbs free energy relationship is applicable. C O R R O S I O NAREA
Figure 2. Freely corroding iron in aqueous HCl.
found as an oxide (ore) in nature. Refining reverses the corrosion process, but yields an unstable product. A note of caution is important here. Although thermodynamics predicts which metals will corrode as an ultimate reality, t i m e is not involved. As an example, diamond is less stable than graphite a t room temperature and atmospheric pressure. This transformation is, however, immeasurably slow and diamond jewelry is safe for many, many generations. Similarly, aluminum is quite thermodynamically unstable a t room temperature and atmospheric conditions, but due to a thin, tightly adhering isomorphic corrosion oxide coating, more extensive corrosion is stifled. Corrosion inhibiting passivation or ennobling changes corrosion potentials as we shall see later. Iron, on the other hand, forms a loose oxide which enables the corrosion process to proceed as predicted. The free energy change of a corrosion system such as oxidation going from one state to another (AG) is a measure of the drive that a reaction has, i.e., how badly an oxide wants to form. This driving force a t constant pressure is a function of three factors. The heat evolved or absorbed is AH (enthalpy). If LWis negative, heat is given off by the system to the surroundings as in the oxidation of iron which is an exothermic reaction. The degree of disorder of the system is A S (entropy). The dissolution of metal usually yields more random, less ordered products; Le., an increase in entropy is evident. The temperature T is expressed in K. Generally an increase in temperature contributes to a more negative second term, hence a more negative AG. If AG is negative the reaction is favored to go spontaneously as written. If it is positive it will not go as written (although the reverse may). If it is zero, the point of equilibrium has been reached. AGcorrosion
= mcorrosion
- TAScorrosion
If standard states are used, the following equations can be written for a metal oxidation reaction. xM(c) + y/202(g)
M,O,(C)
AGO,,, = BAG"f(products) - BAG"f(reactants) where AGOcoris the free energy difference for the corrosion (oxide formation) process. AGOf =IO for any element inlits normal state. Since AGO for each reactant is zero, AGOcorrosion = AGOf of the oxide. Table I indicates AGOf values for some oxides. These thermodynamic values predict that aluminum has a much greater tendency to oxidize than iron, magnesium, or copper. Electrochemical corrosion requires an anode, a cathode, an electrolyte, and the flow of electrons (current) between the anode and cathode. T o better understand this type of corrosion let us look a t a fundamental corrosion reaction, i.e.. iron freely corroding in aqueous hydrochloric acid. Fe
+ 2H+ + 2Cl-s Fez+ + 2C1- + Hz
The overall whole cell reaction consists of two half-cell reactions
-
+
Anode, oxidation, Fe Fez+ 2e loss of e's Cathode, reduction, 2H+ H2f - 2e gain of e's Sum of half-cell Fe 2H+ Fez+ reactions + Hzl
+
-
Standard oxidation potentials, E o +0.44
0.00 -
+O.44 (whole cell emf
mol Table I1 shows a number of standard oxidation potentials. I t has been established by convention that the standard half-cell oxidation (and reduction) potential for the reaction H2 2H+ 2e be given the value of zero. All other redox potentials are measured relative to it, using an apparatus similar to that shown in Figure 3. The values given are equilibrium values for metals in contact with their ions a t concentrations equal to unit activities and a t 25 "C and 1 atm pressure. A t their 1953 meeting, the IUPAC agreed that the reduction potential for a half-cell electrode reaction would be designated as t h e potential (Uhlig, 1963). Note that the reduction potentials would have a sign opposite to those shown under the oxidation potential column in Table 11. The numerical value would be the same. The advantage here is that it corresponds in sign to the polarity of a connected voltmeter. In Figure 3 iron has a negative reduction potential of -0.44. It is also the negative pole of a galvanic cell in which a standard hydrogen electrode is used. Note further that the flow of electrons is from the iron electrode (anode) to the hydrogen electrode
+
Ind. Eng. Chem. Prod. Res. Dev., Vol. 17, No. 2, 1978
Table 11. Emf Series. Standard Electrode Potentials Electromotive series metal
Electrode reaction
Standard oxidation potential,c E". V
Active end Potassium K-+K++e +2'92 Na Na+ + e +2.71 (More anodic) Sodium Mg Mg2+ 2e Magnesium +2.37 A1 A13+ 3e Aluminum +1.66 Zn Zn2+ +2e Zinc +0.76 Cr Cr3+ 3e Chromium +0.71 Iron Fe Fez+ 2e +0.44 Cadmium Cd Cd2+ + 2e +0.40 Ni Ni2+ 2e Nickel +0.25 Tin Sn Sn2+ 2e +0.14 Lead Pb Pb2+ 2e +0.13 Hydrogen H:! -+ 2H+ + 2e 0.00 Cu Cu2+ 2e -0.34 Copper Iron oxidea Fe-0.02 (--0.60) Stainless Fe-Cr-Ni (--0.60) steelb Aluminum Al.Al203 (--0.60) oxiden Silver Ag Ag+ e -0.80 (More cathodic) Platinum Pt Ptz+ 2e -1.2 Inactive end Gold -1.50 Au Au3+ + 3e Iron and aluminum are rendered passive due t o thin oxide films formed. The more noble potential shows this. Stainless steels are similarly passivated by oxide film formations. For reactions reverse of those written, Le., reductions, a sign change is needed (E" = -E"R) where E " R is the standard reduction potential.
---+
+
+
---
+ + + + + + + +
+ +
171
Another basis for thermodynamic predictions of electrochemical corrosion reactions is the equation AGocorrosion= -nFAE", where n is the moles of electrons transferred in the reaction, F is Faraday's constant (96 500 c , Le., equivalent to one mole of electrons), and AEo is the cell potential. A positive cell potential will lead to a negative AGocorrosionor the prediction of a spontaneous corrosion reaction as written. For the above reaction AGO,,, = -2(96 500) (0.44) = -84 920 J. This indicates that iron will corrode in aqueous HC1 with a driving force of -84 920 J. Note that F = 96 500 C = 96 500 VC/V = 96 500 J/V. The equilibrium constant can be calculated from the relationship AGO = -RT In K . Rearranging and substituting yields K = g 1015. This is an expected large equilibrium constant. For potentials a t other concentrations, the Nerst equation can be used: m M n N . . . qQ r R , . , m moles of M n moles of N, etc. At the condition of equilibrium (where equilibrium principles can be applied)
+
+
K= and since AG",,,
-
+
aQq aMm
= -RT In
AG = AGO
+ +
aRr. *
..
aNn
K = - nFAE"
+ R T l n [ [a UM Qm ~ Ua RN' n. . ...]. ]
then
where R is the gas constant and ( a ) is the activity, e.g., a = [MI? and [MI is the molality (mo1/1000 g of water), y is the activity coefficient. AE" is the emf when all reactants and products &rein their standard states, Le., activities equal to unity. AE is the cell emf (Uhlig, 1963). The more familiar form for two reactants and two products is -0.059 (Q)q(R)' n log (M)"(N)" Note in the special case when all concentrations are equal AE = AE". Note further that a t equilibrium concentrations, no net reaction occurs and AE = 0 so that
AE=AE"-
RFFFSFUCE CFLL
Figure 3. Apparatus for measuring corrosion potential, Fe-H+. (cathode).The circuit is completed by the movement of cations to the cathode and anions to the anode through the salt bridge. On this basis, for solid iron (crystalline) immersed in a solution containing Fez+ ions, the potential difference between it and its ions can be measured with a voltmeter connected to an inert reference electrode. This is the equilibrium potential that exists between Fe and Fez+.If the Fe2+ions are present a t 1 g-atom/L (unit activity), T = 25 "C, 1atm pressure, the potential read is the standard half-cell potential.
A E = emf = cell potential = Eocathode- Eoanode = 0.0 - (-0.44) = +0.44 This is the same value as that calculated by the addition of the two half-cell reactions. An electric potential exists between the anode and cathode areas. In real corrosion processes the anodes and cathodes can "switch" periodically so that uniform corrosion occurs. Both oxidation and reduction must occur a t the same rate for the preservation of electrical neutrality of the iron. If either half-reaction stops, so does the overall reaction. This fact can be used to combat corrosion as we shall see later, but for now let us continue with the treatment of ideal cases.
G O = -
-0.059 log K ; n
In linking electrochemical theory with practical observations it can be seen that since AGO = -nFAE", the cell voltage or voltmeter is really a free energy difference meter. As the whole cell reaction proceeds, the voltage difference between electrodes drops since the free energies of the products and reactants approach each other. At equilibrium, both electrodes have the same voltage and no charge transfer occurs (Mahan, 1963). If the standard electrode potentials are listed in decreasing value as shown in Table 11, an electromotiue series is created which has the hydrogen half-cell reaction listed a t a potential of zero. Other reactions are run relative to it under standard conditions. Metals listed a t the top of the series are most reactive, Le., corrode more readily. The lower end consists of "noble" less reactive metals. A metal above another will displace it from a solution containing the lower metal's ions. Iron, for example, will have copper metal plated out on it when placed in a copper sulfate solution. This is an iron corrosion reaction. Metals above hydrogen will displace it from acid solutions. When two metals are coupled, the size of the potential may
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be determined, e.g., (Sn - Fe) = +0.14 - (f0.44) = -0.26. The higher metal is anodic to the lower. Electromotive series for metals in various environments such as sea water are available in the literature. A limitation of the electromotive series is that often it is difficult to get certain metal salts into the needed solution concentrations for comparisons and calculations. The use of a sacrificial anode coupled to a cathode that is to be protected is a well-known means of corrosion prevention based on cathodic protection and these principles. Passivity. If a steel sheet is first dipped in nitric acid and then placed in a copper sulfate solution, the solution will not corrode the steel. Upon rubbing the sheet with a glass rod, copper immediately plates out as could be predicted from Table 11.This phenomenon is a manifestation of passivity or ennobling of the steel (moving down the emf scale from +0.44 to --0.60 V). The nitric acid solution helped in the formation of a thin passive oxide film. The potential and position of the oxygen treated sheet is approximated in Table 11. Of course rupturing the film causes an immediate resumption of the predicted corrosion reaction. Other examples of passivation are the oxide formations of aluminum (A1203), titanium (TiOz), lead in sulfuric acid (below 90%) forming PbS04, and chromium oxides on stainless steels. If stainless steel is starved of oxygen it will corrode. Placing a rubber band around a piece of stainless steel and immersing the banded metal in dilute NaCl will cause visible corrosion (dissolution of the anodic areas under the rubber band in about 2 weeks). It is also believed that stainless steels (la+% Cr) form the less reactive chromium type intermetallic compound FesCr. Supposedly chromium with its five vacancies in its 3d subshell ties up one of iron’s 3d electrons so that the iron atoms then more closely resemble the chromium electronic structure. Stainless steels desperately need oxygen in order to survive the many corrosive environments in which they are used. The relative position of a stainless steel is shown in Table I1 as having an estimated standard oxidation potential of -0.60 V (Slabaugh, 1974). Passivation may therefore be generally defined as a condition in which an ordinarily active metal is made more noble or passive. The resultant action or effect occurs at the metal electrolyte interface. Many times inorganic or organic inhibitors are added to the electrolyte to act as types of passivators. The inhibitors may increase the activation energy for the corrosion process. Overvoltages for the reduction of hydrogen ions create a diffusion barrier on the electrode sites. For every metal-environment corrosion system there is a critical passivation potential called the Flade potential. Inhibitors may increase the rate of metal oxidation to where it passes through the Flade potential and then self-passivation occurs (Powers and Roebuck, 1965). Another means of predicting corrosion reactions is the Pourbaix diagram, which deals with thermodynamic equilibria. The pH of an electrolyte often has a profound effect on the possibilities for corrosion. The Pourbaix diagram shows the possibilities of corrosion considering the pH of the electrolyte (corrodant) vs. the potential difference between the pure metal and its ions in solution. It indicates which pHpotential conditions might lead to corrosion. The generalized Pourbaix diagram, Figure 4, is for the iron-water system. The regions of corrosion, immunity, and passivity are displayed. An important feature of this diagram is that it shows how pH adjustments and/or potential changes (anodic and cathodic) can place normally corroding metal such as iron into “safe” regions. Care must be exercised, however, in employing neutralization (increasing pH) and anodic protection means since moving from the (A) corrosion area upward or to the right increases the propensity for corrosion. Also the diagonal line between the (A) corrosion area and the passivation area represents conditions that could lead to pitting.
+1,
+0,8
0
2
14
6 PH
Figure 4. Generalized Pourbaix diagram for corrosion of metallic iron.
4CT IVATIOH ENERGY
RE4CTION PATt’ PRSD’IC-S
REACTASIS
Figure 5. Generalized driving force-reaction coordinate diagram fur an unimpeded metal corrosion reaction.
u
y” 0
z
T
I
I
c a
EO
v
-
REACTiOil PATt
Figure 6. Generalized driving force-reaction coordinate diagram for typical metal corrosion reaction.
The (E) line shows the reversible potential for the reduction cathodic reaction: 2H+ 2e G H2t. The (F) line shows the reversible potential for the reduction cathodic reaction: 2H+ l/ZO2(g) 2e G HzO(1). Water is stable in the area between the two lines. Below lower line (E), water can decompose yielding gaseous hydrogen; the solution pH would rise. Above upper line (F) water can decompose to yield gaseous oxygen. The solution pH would drop (LaQue and Copson, 1963).
+
+
+
Kinetics The barrier between what is predicted by thermodynamics and what is realized by kinetics is the energy of activation barrier. If as in the so-called ideal case of iron freely corroding in aqueous HC1, the resistance between anodes and cathodes is virtually zero and if reactants can combine and the products leave the reaction site without inhibiting the reaction process, then the corrosion would be essentially instantaneous. Figure 5 shows that situation using a generalized reaction coordinate-potential (energy) diagram. Consider a simple “ideal model” electrochemical cell having
Ind. Eng. Chern. Prod. Res. Dev., Vol. 17, No. 2, 1978
173
w
REACTION PRTH
I
PRO-LC-
RtRiTt.\F
I
Figure 7. Generalized driving force-reaction coordinate diagram for
IC%
IClR
:cc
a corroding metal showing accelerated and inhibited reactions.
I
Figure 9. Generalized mixed polarization (E-log i).
I
CATHODE PO LARIZRTI C:! 2;-
Ze+"*
.
Figure 8. Polarization (E-i) for iron in an aqueous acid solution. one inert electrode, an electrolyte, and a metal oxide forming at the other electrode. The emf (AE)of the cell is therefore proportional to the driving force (energy) of the metal to metal oxide reaction. In Figure 6 the reaction progress (corrosion) is again plotted vs. a driving force. In Figure 5 the reaction is unimpeded. In Figure 6 an activation energy barrier is evident. The reaction is exothermic: the difference between the energies of the products and reactants is the heat of reaction. The activation energy of the reverse reaction is greater than that of the forward. This activation barrier can control the reaction (corrosion) rate since: (1) metal atoms must free themselves from their crystal lattices, which takes energy; (2) oxygen atoms must be made available from oxygen molecules; (3) the production of metal oxide may make the subsequent corrosion reaction more difficult as in the case of corroding aluminum; (4) byproducts collecting a t the reaction sites may further inhibit the accessibilities of reactants. This could manifest itself as polarization. For example, the rate of iron corroding in aqueous HC1 can be decreased by the accumulation of H2 at the cathodic sites. Figure 7 shows three similar corrosion reactions. The inhibited reaction, i.e., higher activation energy, could be due to passivation, inhibitors, or a collection of by-products. Of course the corrosion rate would be lower. The opposite effect is shown in the accelerated reaction curve. Note that the product and reactant energies for the three reactions are the same. Only the rates to the equilibrium point are changed. The Arrhenius equation relates the corrosion rate to the activation energy EAF:k = Ae-Em/RT.The rate constant is k . A is a preexponential or frequency factor. R and T were defined earlier.
I
LOG I
Figure 10. Cathodic polarization (E-log i).
\
~
IC09
LOG I
Figure 11. Anodic polarization (E-log i). The effect of temperature on the reaction rate can be determined from the Arrhenius equation (assuming the same reaction occurs). A and EAFare independent of temperature to a first order of approximation. Since electrochemical corrosion can be represented to a large extent by metals going into solution as cations and since a corresponding number of electrons must flow from the anode to the cathode, measuring the current should yield a corrosion rate. If there is an infinite resistance between the anode and cathode, the potentials of the anode and cathode would remain at their half-cell values; Le., the current would be zero and no corrosion would occur. If, as mentioned earlier, the resistance is zero, the current would be infinite (Ohm's law, E = I R ) . As
Ind. Eng. Chem. Prod. Res. Dev., Vol. 17, No. 2, 1978
174
w
3 4 c E p_
1 INCLLSTC\S
/ WORK P A R D r V I D X 3 r A S
A L L O Y CHRVGES
Figure 15. Possible defects in a metal. I CURRENT
DENSITY,
LOG I
Figure 12. Idealized anodic polarization curve (potential E-current density log i).
Normally the anodes and cathodes are close together so that R is small and for a finite maximum current, E,1 - E,, must be small. The intersection of the anodic and cathodic reaction curves in Figure 8 shows the current value i,. This in turn indicates a corrosion potential of E,,,. The corrosion potential can be measured with an apparatus similar to that shown in Figure 3. In this case the corrosion potential between the metal and its electrolyte is measured and not the cell potential discussed in the earlier section. The corrosion current density can be related to an actual corrosion rate, i.e.
Cmpy= 0.13Ie/d where Cmpyis the corrosion rate in mils penetration per year, I is the current density (pA/cm2, e is the equivalent weight of
CURRENT
DENSITY,
LOG
I
Figure 13. Idealized anodic polarization curve (potential E-current density log i).
the metal, and d is the density of the metal (g/cm3). Polarization curves are used to calculate corrosion rates. The two common methods are by Tafel extrapolation and linear polarization. The slope ( A E l A I )of a potential vs. applied current curve is constant for small (ca. 10 mV) deviations from the corrosion potential. Also (AElAZ)is related to corrosion current lcor
=
[BaBcI [2.3(Ba B,)]
+
where B , and B , are the anodic and cathodic Tafel slopes, respectively. For most linear polarization measurements, B,B,/(B, + B,) is assumed to be constant so that the final equation reduces to i,,, = K ( A I / A E )
CURRENT
DENSITY,
LOG I
Figure 14. Idealized anodic polarization curve (potential E-current density log i).
a practical matter, the resistance approaches zero yielding a finite current. However, when current flows between the anode and cathode, both potentials change (polarize). This polarization results in reducing the potential difference (driving force) between the anode and cathode sites. The flow of electrons between the anode and cathode a t any point can be shown as il =
[ E ~-IEal]/R
where E,l and E,l are the cathode and anode potentials, respectively, and R is the total circuit resistance which includes metal, solution, and films. Figure 8 shows the polarization of iron in an acid solution (Henthorne, 1971b). Fe
+ 2HC1-
FeC12 + H2t
Instruments that measure corrosion rates by linear polarization techniques are available. Polarization Diagrams. Polarization refers to the changes in electrode potentials, during electrochemical corrosion, when electrons flow between the anode and cathode sites. The corrosion (equilibrium) potential and current density can be measured. Possible changes are shown in the idealized Figures 9 to 11. Depolarizations are also possible. Potentiostats. By using an external circuit the potential between a freely corroding metal and its solvated ions can be changed. Changing this potential results in altered degrees of corrosion. A potentiostat is used to hold or change a potential as required. Anodic polarization curves yield valuable information concerning passivity. The idealized Figures 12 and 14 show how changes in anodic polarization can increase possibilities for passivity. In Figure 13 passivity is not likely to be made possible in many environments. In actuality, Figures 12 to 14 would be more curved. Temperature effects would cause shifts to the right, inhibitors/passivators to the left. Nonideal Realities Many nonideal realities plague the corrosion engineer. Metal surfaces, for example, are not well behaved. Consider
Ind, Eng. Chem. Prod. Res. Dev., Vol. 17, No. 2, 1978
h e do not know What we do ~know _ _ _ _ _ _ _ _ _ _ _ - - What --
Electrochemical Theory Presence of anodes and A unified theory of localized cathodes on corrosion. corroding surfaces. Relation between current flow and corrosion. Corrosion potentials of many metals and alloys. Effects of polarization and resistance to current flow Mechanism by which a metal atom leaves its lattice and becomes a hydrated ion in solution. Nature and extent of overvoltage factors. Stress Corrosion Some of the metals and The mechanism of stress environments where corrosion cracking in most systems. stress corrosion can be Role of hydrogen in stress expected, i.e., brass in ammonia, steel in caustic, corrosion cracking. and stainless steels What distinguishes an in chlorides. environment that causes cracking from one that does not. Pitting Electrochemical aspects of pit Why a pit starts at one point propagation. rather than another. The mechanisms by which Physical circumstances that favor pitting. some alloying elements, Why some ions such as such as molybdenum, chloride ions are more improve resistance to active in starting pitting. pitting . Impingement Attack Relative merits of different What determines the ability alloys in resisting of an alloy to form an impingement attack. adherent protective film. Problable mechanisms that are What determines the involved. protective film repair rate. How aluminum improves protective film on brass and iron improves protective film on cupro-nickels.
the many types of corrosion possible by the following “typical” metal surface, Figure 15 (Moher, 1976). If metals were not handled a t all and could be kept in a controlled environment, then corrosion would not be as great a problem. However, even in the manufacture of simple articles, stresses are put into metals. These cause cathode-anode areas that will corrode when a suitable electrolyte is provided. An ordinary iron nail when placed in an aqueous gel containing phenolphthalein and potassium ferricyanide will show a pink (cathodic, due t o OH-) color near the shank and blue color a t the work hardened head and tip (anodic due to iron going into solution). This is just one example. Many more could be cited. The microstructure of metals is undergoing detailed study in attempts t o better correlate microstructural features with corrosion in various environments. A number of other corrosion environment changes can affect corrosion processes. Some are differential solution concentrations, anode-cathode sites, polarization with by-
What we do know
175
What we do not know
-
Cavitation Erosion In a general way, the How to make an accelerated circumstances under test that will rate which cavitation materials properly. damage may occur. Relative merits of different alloys in resisting cavitation erosion. Exact mechanism of cavitation damage. Relative importance to mechanical and chemical factors. How to solve some problems by changing design or by controlling corrosivity of environment, as with inhibitors. Cathodic Protection How to protect metals in some applications, e.g., underground pipes and ship hulls. How to monitor cathodic protection by potential measurements. Mechanism by which cathodic protection is achieved. Significance of potential measurements used to monitor cathodic protection. Atmospheric Corrosion That there is a great spread in corrosivities of atmospheres at different locations and that this spread is due largely to atmospheric pollution. How to measure pollution, humidity, temperature, etc., quantitatively. How to use measurements of pollution and other factors to estimate the probable corrosivity of a particular atmosphere.
products, inhibitors, differential aeration, depolarizations, incomplete passivity, erosion, and cavitation. Corrosion is a complex phenomenon. LaQue (1963) summarized a number of factors that were known as well as unknown about corrosion. This table has been updated. See Table 111. These unknown factors are the basis of ongoing research investigations. Technically, a major aspect that seems missing is “a unified theory of localized corrosion”. From a communications standpoint we still need more effective interactions between the industrial and academic communities. Both have valuable and much needed information, but often a common language is lacking (Donahue, 1977).
Literature Cited Donahue. F., University of Michigan, private communication, Dec 13, 1977 Henthorne, M., Chem. Eng., 127 (May 17, 1971a). Henthorne. M.. Chem. Eng., 99 (June 14, 1971b).
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lnd Week, 43 (July 21, 1975). LaQue. F. L., Mat. Des. Eng , 99 (Jan 1963). LaQue.F. L.. CoDson, H. R.. "Corrosion Resistance of Metals and Allovs". , , 2nd ed, p 89, Reinhold, New York, N.Y., 1963. Mahan, B. H., "Elementary Chemical Thermodynamics", p 102, W. A. Benjamin, New York, N.Y.. 1963. Mazia, J., Met. Finish., 22 (Dec 1977). Moher, J. E., Met. Finish., 73 (May 1976). Parr, J. G., Chem. Eng.. 166 (July 6, 1964). Powers, R, A , , Roebuck, A , ,-,, ,~Encyc,opedia of ChemicalTechnology,t,vel, 6. D 300. 1964. Slabaugh, W. H., J Chem. Educ., 41, 218 (1974).
Uhlig, H. H., "Corrosion and Corrosion Control", pp 23, 28, Wiley, New York, N.Y., 1963.
Received for reuiew October 3,1977 Accepted March 9, 1978 Presented as part of the Symposium on Interfacial Phenomena in Corrosion Protection at the Division of Organic Coatings and Plastics Chemistry, 173rd National Meeting of the American Chemical Society, New Orleans, La., March 1977
Novel Corrosion Resistant Alloys by Ion Implantation V. Ashworth," W. A. Grant,' and R. P. M. Procler Corrosion and Protection Centre, University of Manchester Institute of Science and Technology, Manchester, M60
IQD,England
Since corrosion is essentially a surface phenomenon, surface alloying will often serve to impart adequate corrosion resistance to a metal with intrinsically inferior properties. However, improved corrosion resistance is afforded generally only by alloying additions that remain in solid solution in the base metal. Thus the range of alloying additions that can be used in conventional surface alloying processes to produce corrosion-resistant alloy layers is limited to those which have reasonable solid solubility in the base metal. Ion implantation provides an alternative method for producing corrosion-resistant surface alloys. Furthermore, since the technique requires that the alloying element be ionized and accelerated into the base metal, the limitations on solid solubility imposed by equilibrium phase diagrams are no longer applicable. The paper discusses experimental results showing how the ease of passivation and resistance to environments containing chloride ions of iron, aluminum, and a type 304 stainless steel can be improved by producing novel surface alloys by ion implantation. Finally, some results are presented to show how ion implantation may be used as a research tool in a corrosion investigation.
Metallic corrosion has been defined as the passage of the metal into the chemically combined state (Shreir, 1963).It has been shown (Evans and Hoar, 1932) that in aqueous solutions the reaction occurs through an electrochemical mechanism in which the metal is oxidized (loses electrons in an anodic reaction) while solution species are reduced (gaining electrons in a cathodic reaction). Clearly the sites of these anodic and cathodic reactions must be linked by electronic and electrolytic pathways in the metal and the corrosive environment, respectively, if the requisite charge transfer is to occur. Further, it is a matter of experience that when corrosion takes place there is no discernible buildup of charge, from which it may be inferred that the rates of electron release (anodic process) and consumption (cathodic process) are equal. The existing methods of corrosion prevention and control are necessarily based on attempts to interfere with these prerequisites for corrosion. Thus corrosion control techniques fall into five broad categories, viz. materials selection, environmental control, electrochemical protection, the use of coatings, and design. Careful selection of materials may minimize corrosion, since for reasons that are readily explained by recourse to electrochemical thermodynamics and kinetics, not all metallic materials are equally susceptible to corrosion in a given environment. By contrast, it may be possible to modify the corrosive environment either by removing the agressive agent entirely or by making judicious additions that in the presence of the aggressive agent interfere with the electrochemical kinetics of oxidation or reduction and 1 Department of Electrical Engineering, University of Salford, Salford, England.
0019-7890/78/1217-0176$0.100/0
thereby inhibit the corrosion reaction. Alternatively, it may be possible to intervene in the corrosion reaction electrochemically, to stimulate either the anodic reaction, thereby promoting protective film formation (so-called anodic protection), or the cathodic reaction (cathodic protection) with a corresponding lowering of the metal oxidation rate. A too often neglected method of controlling corrosion is by means of good design. Bad design can completely negate subsequent good materials selection or protection schemes by fostering avoidable corrosion cells. Good design can permit the use of less expensive materials and protection schemes. Finally, corrosion may be controlled by the application of coatings that inter alia serve to separate the metal from the corrosive environment and may also have other beneficial effects. On a tonnage basis, organic coatings (paints, polymers etc.) protect more metal from corrosion than any other type. Inorganic coatings are less commonplace but find increasing application. As a class they range from the anodic oxide films produced on aluminum and the phosphate chemical conversion coatings on steel to the applied coatings based on vitreous enamels, glass, and ceramics. Both these coating types are effective largely because they reduce the interfacial conductivity and thereby provide a high resistance path between the anodic and cathodic reaction sites. Metallic coatings are somewhat different, and their use may be regarded as an exercise in metal substitution; Le., the coating presents to the environment a more corrosion resistant material than does the substrate itself. Metallic coatings that are less noble than the substrate material provide the added advantage that, if damaged, they confer temporary electrochemical protection on the substrate pending reassertion of
0 1978 American Chemical Society
