WILSON L. ORRLand HENRY A. STAFFORD University of Southern California, Aeronautics, Santa Maria, California
School
of
M o m textbooks on chemistry and metallurgy are too An anodic and cathodic reaction is always necessary brief on the theory of corrosion while the more complete if corrosion is to take place since the metal cannot lose works on corrosion are too lengthy for the time most electrons unless some other substance can accept them. students can allow to the subject. The purpose of this If the metal is in an acid solution the cathodic reaction paper is to give engineering students a brief survey of the is usually theory of corrosion and its application intermediate 2H20f 2 electrons = Hz + 2H20 (Reaction 2 ) between the two extremes. This paper is necessarily The symbol HsO+ is used here as anhydrated ion brief and gives only one or two examples to illustrate each point but should he sufficient to give the student because the proton, H + exists. Probably the most acthe necessary background to apply the theory and see curate representation would he H(H20).+, but H30+ and simply H + are the most frequently used symbols. other examples as he encounters them in his work. The importance of the prevention of corrosion is evident since it is a major cause of structural failures in TABLE 1 bridges, machines, aircraft, automobiles, etc. Its cost Electromotive Series yearly is tremendous not to consider the loss of lives and injuries caused by such failures. I t is evident that potenpotenif engineers are to design and maintain structures Reaction tial Reaction tial properly they must keep corrosion in mind. An under= Lit + e3.02 = Co++ 2e0.28 standing of the fundamentals of corrosion and the 2.92 = K* e= NiC+ + 2e0.25 factors influencing it are necessary if failures due to = Sn++ + 2e2.87 = C a t + +2e0.11 = Naf + e- 22.34 = P h + + + 2e0.13 .il corrosion are to be kept at a minimum. = Mg++ + 2e= 2H+ + 2e000 One could hardly seriously ask the question, "Why = BeC+ + 2e= C u + + + 2e1.70 -0.34 do metals corrodef" since the occurrence of metals in 1.67 = Al+++ + 3e= Fef*+ e-0.77 = M n f + + 2e1.05 = Ag+ e- -0.80 nature as compounds rather than pure metals in all = Zn++ + 2e= HgC+ + 2eOi6 -0.85 cases except the noble metals indicates that the com= Cr+++ + 3e0.71 = Pt++ 2e- -1.2 = Fe++ + 28= OZ4H++4e0.44 hined form is the more stable. I t is rather more surprising that metals corrode as slo~vlyas they do. Cor- Cd = Cd++ + ze- o,,o A4u = Au+ + e - -1.23 rosion then naturally takes place when the environment I t should be noted that the potential values given in this tahle the a chance to combine with other apply only to the condition where the metal is in contact with ments. a solution in which the activity (activity means effective can-
+
+
+
+ +
ELECTROCHEMICAL THEORY O F CORROSION
Fundamentally the rusting or corrosion process is one in which electrons are lost by metallic atoms vvhichthen become positive ions. Fe = Fe++
+ 2 electrons
~h~ area where the metallic atoms lose electrons and become ions is called the anodic area and this reaction the anodic The area where the electrons are used is another chemical reaction called the cathodic area and this reaction the cathodic reaction. Since metals are good conductors of electrons these areas may be separated if environmental conditions warrant such.2 Present address, Hanoock Research Foundation, University of Southern California, Los Angeles, California. 2All corrosion is essentially electrochemical in nature since electrons are always transferred from the metal to same reducible substance but some prefer t o view corrosion as eleotroohemical only when the anodic and cathodic areas are separated by a discernible distance. If the corrosion is uniform over the entire surface they prefer to call it direct chemical attack.
centration) of the ion indicated is one mole per 1000 g. of water. I n any other solution, different values for the potential \\-auld develop. The student is referred to any general or physical chemisbry textbook for details on how the values of the oxidahion potential are measured. Notice the oxiddon potentid of H2 is arbitrarily taken as zero and all others are measured in reference to hydrogen. (Example) The listed oxidation potential for zinc is 0.76 which means that the reaction, Zn = Zn++ 2e-, tends to proceed toward the right (oxidation) with sufficient energy to drive the reaction, H2 =. 2H+ 2e-, toward the leit (reduction) with a net electromotive force of 0.76 volt for the combined reaction, where zinc is in contact with a solution of unit activity (effective oonoentrstion of one mole per 1000 g, of water) in z n + + , and HI gas a t one atmosphere pressure in contact with a. solution of unitactivityin H+.
+
+
However, in engineering we are usually more interested in atmospheric corrosion where the usual cathodic reaction is O1
+ 2H20 + 4 electrons = 40H-
(Reaction 3 )
It k necessary that charged Or groups of atoms as Fe++ and OH- migrate from the
APRIL, 1950
position where they are formed and pair off with ions of the opposite charge so there is no great accumulation of charge in any area. Ions migrate rapidly and easily in solution hut much slower along the surface of dry metals. Therefore metals in solutions may develop anodic and cathodic areas quite far apart hut in the atmosphere they will usually be closer together. Metals differ considerably in their tendency to lose electrons and become ions (oxidation potential). The electromotive series places metals in a series according to their tendency to lose electrons. Those having the greatest tendency to lose electrons (highest oxidation potential) are a t the top of the series and those with little tendency at the bottom. Then the higher a metal in the electromotive series the greater is its tendency to lose electrons, i. e., to become oxidized. Oxidation in this sense is not limited to union with oxygen hut refers to the process of losing electrons whether to oxygen or any other substance. The position of a metal in the electromotive series (Table I), while very useful, cannot be accepted as an absolute indication of its activity in reference to other metals under all conditions because the environment is also a determining factor and the laboratory measure ments of oxidation potentials are under different conditions from those encountered in the actual use of the metal. The galvanic series (Tahle 4) is of more value, as far as corrosion is concerned, than the electromotive series. The galvanic series has basic characteristics which make it analogous to the electromotive series but it takes into consideration over-all and practical aspects. It is based on experience in corrosion testing under actual operating conditions in numerous corrosives and on practical results with metals and alloys in service. However, since different environments cause the potential to shift, this series will not allow absolute predictions of galvanic action in all corrosive environments. There are other factors such as crystal size, strains, t v ~ eof surface, surface films of corrosion vroducts. etc., that tend to either raise or lower the potential: When any factor causes a metal to behave electrochemically as though it should be lower in the galvanic series than it actually is, it is said to he in a passive condition. Metallurgists often put metals in the series for hoth active and passive condition where the passivity is established by certain treatments or specific factors of the environment so the position can he definitely established. Notice in the galvanic series (Tahle 2) the chromium-iron, nickel, and 1&8 chromiumnickel-iron alloys are listed for hoth active and passive condition. In general the rate of corrosion will he greater, the greater the potential difference in the various areas of the metal as caused by strain, surface coating, crystal structure, inclusions, contact with another metal or differences in environment in different areas. Oxidation (anodic reaction) will take place where the potential is the highest and reduction (cathodic reaction) at the centers of lowest potential.
203
TABLE 2. Galvanic Series of Metals and Allow Corroded end (anodic or least noble) Magnesium Magnesium alloys Zinc Aluminum 25 Cadmium Aluminum 17ST Iron or Steel Cast Iron Chromium-iron (active) 1- chromium-nickel-iron (active) Lead-tin solders Lead Tin Nickel (active) Inconel (active) Brasses Bronzes Copper-nickel alloys Nickel ( assive) Inconel ppassive) Chromium-iron (passive) 18-8 chrominm-nickel-iron (passive) Silver Gold Platinum Protected end (cathodic, or most noble)
FACTORS INFXUENCING CORROSION
By use of the ferroxyl indicator8 (5)me can actually see where these areas develop and therefore observe how various conditions, such as-strain, cause corrosion to be accelerated in iron and steel. The ferroxyl indicator is an agar-agar gel containing potassium ferricyanide and phen~lphthalein.~The ferricyanide ion reacts with the Fe++ ions as they are formed and migrate into the gel forming a blue color due to Fea[Fe(CN)a]z and the phenolphthalein forms ared color in the presence of hydroxyl ions formed by the cathodic reaction. If we place an ordinary iron nail in this gel and allow it to stand a few hours we can observe the development of the anodic and cathodic areas as shown in Figure 1.
- - -- F i g " . . 1. Iron Nail in Parroxgl Gel Diagonal lines indicate m o d e (blue): horiaontal lines indiaate cstbode (red).
The position of the blue color indicates very clearly that iron in a strained condition is more susceptible to corrosion (that is more anodic) since its formation on the nail iis a t the spots where Strain would have resulted from the forming of the nail. To demonstrate this A similar reagent for demonstrating the corrosion of aluminum snd aluminum alloys can be used (7). ' Althoueh Meldrum ( 5.) does not mention the use of NaCl in the rcnpel.t, it k d t m added since i t increases the rate of corrosion and tl.rrefore r1.r rolored r r m r Hpprnr in a shortcr rime.
.
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204
further, if we place a bent nail in this gel we observe metal. If we take a piece of polished iron or steel which results as shown in Figure 2. Here again we observe is free from any strain and partially immerse it in salt that the strained areas are more anodic than the un- water, corrosion proceeds much faster than if it is strained areas. completely immersed. On this uniform surface all the metal should have the same tendency to lose electrons, but a t the air-water junction there is a much higher concentration of oxygen and water which is necessary to remove electrons by Reaction 3; therefore this reaction has the greatest tendency to take place a t this point. Another factor that makes the reaction proceed is the ease with which Fe++ ions can be removed as formed from the submerged part of the iron because the negative C1- ions exert an electrostatic force on them Figure 2. Bent Iron Nail and help pull them into the solution. I t is not necesDiagonal lines indioate anode sary to have an electrolyte in the water although it (blue); horizontal lines indieete greatly speeds up the corrosion. Even in pure 11-ater, cathode (red). The position of the red areas (cathodic areas) will in general be a t the points having the lowest oxidation potential, i. e., the least tendency to lose electrons. In Figures 1 and 2 this is the unstrained iron. The cathodic areas are usually larger than the anodic areas. If other materials of very low potential, such as slag inclusions, surface films, or less active metals, are present they will be the centers of the cathodic reaction. We may demonstrate this by attaching a less active metal, copper, to iron and observe its corrosion in the ferroxyl gel (Figure 3). The metal of lowest oxidation potential
-------=- -4-5 -
\\\\''
"""'
Fi3. Iron Neil Wound with Copper Wire Diagonal lines indioate anode (blue): horisontsl lines indicate oathode (red).
(copper) becomes the cathodic area and the iron the anodic area. If we place zinc in contact with iron we observe the same type of behavior except in this case the iron is the metal of lower potential and therefore becomes the cathodic area, while the zinc undergoes the anodic reaction Zn = Zn++ + 2e-
(Reaction 4)
A white area composed of zinc ferricyanide develops around the zinc (Figure 4). 3Zn++
+ 2 F~(CN)B---= Zna[Fe(CN)&
F 4 Iron Nail through Zinc in Femoxyl Csll Diagonal Lines indicate anode (white): horirontal lines indicate eathode (red).
due to the polar nature of water, the ferrous ions can go into and move through the 'solution much faster than along the surface of the metal in air. At the water-air junction hydroxyl ions are formed by Reaction 3, and in the submersed :lortion of the iron ferrous ions are formed and go inio solution. Some of these Fe++ are oxidized to F;+++ ions by the oxygen. A precipitate of Fe(OH)%and Fe(OH)3forms when the concentration of the ions exceeds the solubility product constant for Fe(OH)% and Fe(OH)3. The Fe(OH), in this precipitate may then be further oxidized to Fe(OHj3. Whether the initial precipitate is largely Fe(OH), or Fe(OH)3 depends on the relative concentrations of Fe++ and Fe++ ions. Most of this "rust" will form in the solution but that which does form on the metal will accumulate at the boundary between the anodic and cathodic areas. Complete equations for rust formation from iron are given below. +
+ + + ++
Fe++ 20H- = F~(OH)E and 0 s = ~ e + + + 4 OH4Fe++ 2Hz0 4Fe(OH)r On 2HsO = 4Fe(OH)>and Fef++ 30H- = FmH)a -
+
+
If the rate of corrosion is observed by the time required for these colors to develop it will be noted that Notice if 2Fe(OH)3 is written as a hydrated oxide the greater the difference in potential of the anodic instead of a hydroxide it is FezOs.3Hz0 which is the and cathodic material the more rapid the corrosion. formula usually given for rust. When the potential difference is between two metals, as Localized anodic and cathodic areas will not exist if in the two examples just cited, the corrosion is known the metal is pure, smooth, free from strain, and exposed as galvanic action. to a uniform environment. In this case both reactions Differences in potential can also be set up by differ- (3) and (2) should take place uniformly over the entire ences in concentration of substances in contact with the surface producing a uniform film of corrosion product.
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This uniform film offers considerable protection to the underlying metal by separating i t from the environment. The extent to which such a uniform film protects the metal depends on how firmly it adheres to the metal, i. e., how easily it is chipped and flakes off. If the film is chipped or flakes off then the difference in potential between the oxide film and the exposed free metal are sufficient to accelerate corrosion. If the exposed anodic area is much smaller than the cathodic area (which is usually the case) then the type of corrosion known as pitting develops. The iron in these small anodic areas continues to form ferrous ions and the rust accumulates a t the point where the ferrous ions meet the hydroxyl ions formed by the cathodic reaction. Thus the small anodic area becomes deeper and deeper -a pit. Since pitting is the most destructive type of corrosion, let us examine it more thoroughly and see why the pit continues to corrode deeper. On the surface there are usually some points that are deeper than the main surface, perhaps from a scratch, air bubble, or irregularities from some other cause. If any of these irregularities resulted in straining the metal, the anodic reaction will tend to take place in the strained metal. Whether or not the indentation is strained, the supply of oxygen will be slightly smaller a t the bottom of a pit than a t the high points of the surface. The location of the highest concentration of oxygen will be the cathodic area since oxygen is the determining factor in the rate a t which electrons can be removed (Reaction 3), and the location of the lower concentration of oxygen will he the location of the lowest cathodic activity, or the highest anodic activity. This factor is always acting in conjunction with the stress concentrating effect of the pit. The structural properties of notch-sensitive materials are rapidly reduced in this manner, i. e., aluminum alloys are more seriously affected than steel. That; relatively small differences in concentration of oxygen can determine the location of the anodic and cathodic areas, can he shown by placing an iron nail in a test tube of ferroxyl gel and after several days observing the result as shown in Figure 5. The top end of the nail, being closer to the air, will have more oxygen diffusing through the gel to it than the lower portion. The red area (cathodic area) develops near the ton and the anodic area near the bottom. Notice that &is differs from Figure 1 where the nail was covered by a thin layer of gel in a flat dish in which the amount of dissolved oxygen is practically uniform over all the nail. The recent observation (11)that freshly precipitated Fe(OH), adsorbs oxygen to a considerable extent seems to explain even better why rust spots become centers of cathodic activity, especially if the Fe(0H)a retains the ability to adsorb oxygen for some time. Adsorbed oxygen on the rust would greatly increase the tendency of the cathodic reaction to take place on the rust and in this way speed up the rate a t which the iron next to the rust would lose its electrons--or corrode. This explains very well why the iron tends to corrode and pit
205
a t the edge of rust spots and gradually work under them. Since corrosion is accelerated by diierences in oxidation potential in different areas we would expect an alloy to he more resistant to corrosion if it is a singlephase alloy (a solid solution) than if it consists of crystals of one composition dispersed in a matrix of a different composition. This expectation is generally verified. Heat treatment of alloys is the operation which determines the distribution of the constituents of the alloy. Therefore the exact heat treatment is often of great importance in determining the rate of corrosion for a given alloy. The Cu-A1 alloys such as 17s-T and 24s-T which contain about 4 per cent Cu offer a good example of the importance of heat treatment on the corrosion resistance of an alloy. This amount of copper is completely soluble in aluminum at 900°F. but is not entirely soluble a t room temperature. It is common practice to heat treat a part made from one of these alloys by heating it to about 920°F. to effectsolid solution of the copper in the aluminum and then rapidly quenching it in cold water to retain the copper in solid solution. On aging for about four days the hardness and strength of the alloy gradually increase. This is believed to he due to CuAla precipitating from the Cu-A1 solid solution. However, these Cu& particles, if present, are in a very finely divided or critical state and the alloy, as far as corrosion is concerned, behaves as if it were a single-phase alloy. If the rate of cooling from 920°F. is slower, the CuAll crystals have time to grow larger and the portion of the solid solution adjacent to the CuAb crystals becomes depleted in copper and therefore relatively purer
Figure 5. Iron Nail in Felloxyl 0.1 Diagonal lines indicate anode (blue); harirantal lines indicate cathode (red).
aluminum which is anodic to both the main mass of Cu-A1 solid solutiou and the CuA12. This condition leads to a type of corrosion known as intergranular corrosion since this anodic zone will he corroded out following close to the grain boundaries. This is one of the most dangerous types of corrosion because it is not easily detected until the part fails (4). METHODS OF CONTROLLING CORROSION
There are many ways of controlling corrosion and the
206
JOURNAL OF CHEMICAL EDUCATION
method which is best for any given use depends on many factors. If the fundamental nature of the corrosion process is carefully considered and examined in reference to the conditions encountered in a given case then an intelligent choice can be made. However, many surprises are found in corrosion in unique environments which simply emphasizes the fact that our theories are still incomplete. Actual test under exposure conditions is the only absolute method of evaluating a given method of corrosion prevention because of the many variables encountered. However, since tests under identical conditions to the purposed use of the metal may take years for accurate evaluation, accelerated tests are usually used. The student is referred to the Corrosion Handbook (9)for information on corrosion testing. Let us recall the two conditions necessary for corrosion to take place: (1) Electrons that tend to be lost by the metal in the anodic reaction cannot be lost unless they can he used in some other chemical reaction which we call the cathodic reaction. (2) The ions formed by both the anodic and cathodic reactions must he removed from where they are formed or must combine with ions of opposite charge so no great accumulation of charge occurs in any area. Any condition which hinders either of these processes will inhibit corrosion, the most important of these are discussed individually in the following paragraphs. Surface Coatings Which Mechanically Isolate the Metal from the Corrosive Environment. Coatings such as paints, lacquers, varnishes, resins, rubber, plastics, corrosion product films, and less active metals than the protected stock, belong to this class. All of the above have one property in common; they protect the metal only by isolat.ing it from materials which can cause corrosion (i.e., reducible substances that can accept electrons from the metal). If the coating is successful in keeping water, oxygen, and other reducible substances from contact,ingthe metal then of
CORROSION PRODUCT (FILM M X ) rimre8. C.,rr..i.,n Pro..edin. through a Corraion Roduct Film Anodic reaction M collosion product M*
- M+
+ X-
+ e:
cathodic reaction X nqx.
+ e-
-
X-:
course no corrosion can occur. In actual practice too little care is exercised in applying these various treatments and they are therefore often not completely impervious to water and air. In service they are often damaged and must be reserviced often to maintain their effectiveness. The greatest disadvantage in using one of the above is that when they are damsged or hroken they may actually accelerate localized corrosion since the exposed anodic area will usually be small. And in the case of the less active metal coatings, the differencein potential between the cathodic coating and the exposed anodic area will accelerate corrosion. The nature of corrosion products in some cases is such that after they are formed they offer some protection by forming a physical barrier between the metal and its environment. If the corrosion product forms a continuous film that is nonporous and prevents water and oxygen from penetrating to the metal then it will stop corrosion. Aluminum is the outstanding example of this. From the position of aluminum in the electromotive and galvanic series (Tables 1 and 2) one would expect it to corrode rapidly. It does in fact corrode very rapidly until a thin film of aluminum oxide is formed which then inhibits further corrosion by isolating the metal from its environment. I t is easy to see why such a film v,ould inhibit corrosion since the onlv through " wav " corrosion could ~roceed a continuous nonporous film would be for the reduciie substances (oxygen, water, or hydrogen ions) to diffuse through the film to the metal or for the film to conduct the electrons from the metal to the reducible substances and allow the metal ions formed to migrate outward to meet the anions formed by the cathodic reaction. This is represented diagrammatically in Figure 6. In most films these are very unlikely processes and do not occur to any appreciable extent, except a t high temperature (6). A film of aluminum oxide, better' than the natural film, can be obtained by making aluminum the anode in an electrolysis cell and using sulfuric acid or a chromic acid solution as the electrolyte. Aluminum so treated is said to be anodized. After anodizing, the film is usually sealed by immersion in hot water which hydrates the oxide, decreasing its porosity. This treatment forms a good basis for primer and paint, but where the structure is subject to vibration or stress it cannot he relied upon without painting. Oils and greases also tend to prevent corrosion by excluding water and oxygen from the metal on which they are used. Cathodic Protection. Cathodic protection refers to any method of protecting a metal by causing it to always be cathodic. This may be accomplished in two ways (1) by attaching the metal to the negative terminal of some source of direct current, and (2) by attaching to it a metal which is more anodic than the local anodic areas in the metal to be protected.
207
APRIL, 1950
Galvanized iron is a familiar example of this type of protection. The iron is covered with a thin coating of zinc. Zinc is anodic to iron6 and therefore undergoes the anodic reaction Zn = Zn++
+ 2e-
and the electrons flow to the iron where the cathodic reaction takes place. This type of coating is often called a sacrificial coating because it protects the main stock by itself being corroded. No rusting of the iron occurs until most of the zinc has first been corroded in the local area. The actual distance of exposed iron 'from zinc which may exist without being corroded varies with the environment. In sea water it may be as much as several inches but in some fresh mraters of low conductivity as small as 0.1 inch (10). Cathodic protection also operates on aluminum alloY% e. g., so-called clad alloys. Aluminum alloys usually corrode faster than pure aluminum but because of mechanical properties it is seldom possible to use Pure aluminum for structural Purposes. An alloy having the desired physical properties may be coated with a thin layer of pure aluminum, or an anodic alloy, and the desired strength and physical properties are retained and the coating gives the necessary resistance to corrosion. The coating material is chosen so as to be anodic to the main stock. where clad sheets are riveted or welded to unclad alloys t h e coating on the clad sheets tends to Protect electrochemically the exposed areas of the unclad alloy. However, for severe atmospheric conditions and where maximum strength is very important as in the aircraft industry, the unclad alloys are always anodized and painted. Another example of cathodic protection used in the aircraft industry is the cadmium plating of steel parts that are to be in contact with aluminum. Cadmium is anodic to iron (Table 4) and thus protects it, and cadmium and aluminum are VeV similar in potential so very little galvanic action results from the cadmiumaluminum contact. However as a further protection the aluminum is usually coated with a zinc chromate primer prior to assembly. Magnesium anodes are being extensively used to Protect iron structures, especially buried pipe lines. Magnesium in the form of a cable or cast blocks or rods is attached to the iron a t regular intervals, the distance apart depending on the conductivity of the medium in which the pipe is buried. Inhibitors and Passivators. The more common inhibitors are organic compounds containing a polar group which causes the molecules to be adsorbed in a monomolecular layer on the metal surface and thus diminish the rate a t which reducible substances such as oxygen, water, and hydrogen ions Can reach the metal.
They are usually used in liquids which are in contact with metals (such as antifreeze). Passivators make up a class of inhibitors, oxidizing agents (i. e., strong electron absorbers), which also form mono-layers on the metal but are more strongly adsorbed than the inhibitors. The bonding of passivator to metal appears to he chemical in naturebut no stoichiometric compound forms and the metal lattice is left intact, ~h~ surface atomsof the metal no longer have the same tendency to lose electrons, i. e., their oxidation potential is lowered, ~h~ ion is the commonly used pltssivator, zinc pajnt so widely as a primer thus actsnot only as p ~ nbut t also as a passivator,. ~ ~ p u r ~t e a t i~o n and l ~ l l ~ ~h~~~ ~ i are ~ ~some . cases where minor impurities in a metal are largely the causes of its rapid corrosion. This is probably simply galvanic action. Iron in magnesium is a striking of this. 1fthe iron of magnesium to is below 0.016 per cent, magnesium is quite corrosion in a 3 per cent sodium chloride solution but if the iron content increases to 0.020 per cent corrosion ~roceedsa t least eighty times as fast (3). Obviously in this case a high degree of purification will greatly decrease the rate of corrosion, 1, other cases alloying a given metal with another metal may produce a metal of better corrosion resist~ h , steels are probably the most spectacular example. They are .iron base alloys containing from 12 to 18 per cent chromium and often other elements among which nickel, manganese, molybdenum, titanium, and columbium are the most mon. They are resistant to corrosion in many media which rapidly attack mostother metals, to ~ h passivity , of stainless steel is usually film of Cr20s; a thin coating of a highly however, recent studies (8) have shown that in 18-88 stainless steel the metal is passive largely because of physically adsorbed gases on its surface. A sample of 18-85 stainless steel loses its passivity when exposed to a reduced pressure (vacuum) a t room temperature and becomes passive again when exposed to air. This process is reversible and the metal can be passivated, activated, passivated, etc., by alternate exposure to air and vacuum at room temperature. hi^ ready breakdown in vacuum at room temperature indicates a weak bond between the gas and the metal. or even chemisorbed gases would he considerably harder to ,move d, probably requirean elevated tern-
,,,,,
perature.
~ l t h o u g hthe mture of the bond between metals in alloys is little understood, it appears that the electron strUctum of the dloy determines the type of film which may form on its surface and that the sharing of electrons between iron and chromium in the stainless steels 6 ~h~~~ is some evidence that ainc either becomes much leas protective or may even reverse its potential with respect to iron acts in such a manner as to allow chemisorption of in hot fresh water. For this particular application then galvsr oxygen other compounds on the surface which nbed iron is inferior to ungalvanieed iron (8). serve, in turn, to satisfy surface valence forces. ChroHot water pipes in Baltimore after several years service showed galvanieed iron pipes to be pitted much more than ungalvaniaed mium possesses this property of itself, hut iron does not, so that when the two are alloyed, their interaction pipes (1).
JOURNAL OF CHEMICAL EDUCATION
203
produces essentially a metal having the same passive properties as chromium, when their ratio is such as to produce the necessary electron structure. Experimentally this is found to be about 12 to 18per cent Cr. Alteration of the Environment. By removing dissolved oxygen and carbon dioxide from boiler feed water the rate of corrosion in boilencan be greatly reduced. It is now common practice to deaerate feed waters for boilen. For high pressure steam plants, the oxygen concentration is usually reduced to less than 0.005 p. p. m. Modern deaerators can accomplish this degree of deaeration hut sodium sulfite can be added to remove the last trace of oxygen. For many industrial systems the dissolved oxygen is reduced to less than 3 p. p. m. by spraying the water into an evacuated chamber where the dissolved gases are effectively removed without heating the water. Carbon dioxide causes corrosion especially in condenser systems by acting as a weak acid. It is removed in the deaeration process.
In general the removal of anyreducible substance ( i . e., oxidizing agent) from the environment of the metal will retard corrosion by limiting the cathodic reaction. REFERENCES (1) BONILLA, C. F., Trans. Electrochem.Soe., 87,237(1945). (2) FONTANA, M.G.,Chern. Eng. News, 25,193A (1947). W. S., in H. H. UHLIG'S "Corrosion Handbook," (3) LOOSE, New York, John Wiley & Sons, 1948,p. 237. R. B.,in H. H. UELIG'S"Corrosion Handbook," (4) MEARS, New York, John Wiley & Sons, 1948,p. 54. W. B., J. CHEM.EDUC., 25, 254 (1948). (5) MELDRUM, (6) MILEY,H. A., in H. H.UHLIG'S''Corro~ionHandbook," New York, John Wiley & Sons, 1948,pp. 11-20. (7) ORE,W.L.,J. CHEM.EDUC., 26, 267 (1949). (8) SCRIKORR, G., Trans. Eledrochem.Soc.,76,247(1939). H. H. (Editor), "Corrosion Handbook," New York, (9) UHLIG, John Wilev & Sons. 1948. . -~ (10) UALIG, H. k,~hem.'Eng.Nezus, 24,3154(1946). K.,AND J. KROLL, Gas u. Wasswfach, 89, 243 (11) WICKERT, ~
