CORROSION, the BILLIONDOLLAR THIEF I.
I n t r o d u c t i o n , D e f i n i t i o n , H i s t o r y , and E l e m e n t a r y C o n c e p t s
FREDERICK A. ROHRMAN Michigan College of Mining and Technology, Houghton, Michigan
This paper, the first part of a three-part series on corrosion, has for its purpose the initiation of the topic to the general reader. The economic importance, history, definition, and elementary fundamentals of the subject are presented with the uiew of acquainting the reader with the most accurate knowledge of corrosion. All corrosion i s shown to be the resultof cell action, i. e., the action of two different elements in the presence qf a n electrolyte. The importance of the Nernst equation and hydrogen ouemoltage i s stressed. The recent work of P a l m e r on "induction period" i s outlined and its importance demonstrated.
losses. A cnrsory examination of our "metal age" shows only too clearly what a staggering problem corrosion presents us a t every turn. Additional evidence on the importance of the problem can he visualized when one realizes that hundreds of specialized experts are devoting all their time to the many phases of corrosion. In this country and abroad specialized committees, organizations, and periodicals are established for the study of its problems. Up to 1932, sixteen books, several bibliographies, and more than five thousand papers had been published upon this subject. DEFINITION
I f gold ruste, What shall iren do? Chaucer
T
HE scientist and the engineer are constantly called upon to direct their attentions to the elimination of waste. As President Hoover has said,' "It is only through the elimination of waste and the increase in our national efficiency that we can hope to lower the cost of living, on the one hand, and raise our standards of living, on the other." Today, the greatest source of waste is corrosion. It is the purpose of this paper to present some of the problems of corrosion and some of the information that has been accumulated on the subject. One phase of the importance of the problem has been emphasized by Sir Robert Hadfield, the great English metallurgist, who stated that the world's debit to corrosion in 1920 was about $2,500,000,000.2 It has also been computed that 33% of all the steel and irou fabricated is for the purpose of replacing units destroyed or damaged by corrosion.3 Bancroft4 stated that 10% of the capital invested in the American petroleum works is charged against the ravages of corrosion. Kendall and Spellerqnd that the cost of gasoline is one cent higher because of corrosion. One of the world's largest manufacturers of chemicals charged $2,500,000 in 1930 to corrosion Them. & Met.Eng., 32, 151 (1925). HADPIELD, Proc. Royal Soc., 101A, 472 (1922). "ALD, "Werkstoffe und Korrasion," Spalrner, Leipzig, 1931, Vol. 1. p. 124. ' BANCROET, Chem. & Met.Eng., 32,6,336(1925). KENDALL AND SPELLER, Ind. Eng. C h m . , 23,740 (1931).
The word corrosion is derived from the Latin rodere (to gnaw). The Latin for rust, however, is robigo or rubigo (red) ; thus, iron rust was called ferrugo; copper rust, aerugo, etc. Speller6 defines corrosion as "the chemical action of certain external agencies on metals which causes their deterioration or destruction." The Deutscher Gesellschaft fiir Metallkunde, the Verein Deutscher Ingenieure, and the Verein Deutscher Eisenhiittenleute have agreed upon a definition of corrosion which states " . . . die Zerstdrung eines festen Korpers, die durch unbeabsichtigte chemische oder elektrochemische Angriffe won der OberfEiiche ausgeht."' This might be translated as, ". . . the destruction of a solid body which follows an undesired chemical or electrochemical attack a t the surface." It will he noted that the German definition eliminates the application of the term to anode corrosion, which certainly is a desired result in some electrochemical processes. Since the Germans define their terms thus, there can be no confusion between desired and undesired corrosion; to them all corrosion is destructive economically. In this paper, however, Speller's definition suffices. HISTORY
The history of corrosion, no doubt, has its origin beyond the records of mankind. Historical references to the metals show that the ancients were familiar with gold, silver, copper, lead, iron, tin, and their alloys, such as electrum, an alloy of gold and silver. It is
-
SPELLER, "Corrosion, causes and prevention," McGrawH i 1 Book Co., New York City, 1926,p. 6. ' KROHNR~. MAAS, AND BECK, "Die Komosion," Hinel, Leipzig, 1929, vol. 1, p. 3.
I.41
difficult to ascertain which of these metals was first used. Quite probably, however, since gold, silver, and copper exist free in nature, they were the first to come into service. Meteoritic iron was also known and used. It has been frequently reasoned that because of the absence of iron specimens among archeological excavations, iron was not used appreciably in the early times. The fact is, however, that iron is so readily corroded that it could hardly be expected to endure the elements for centuries. Hadfield is one of the many who believe that iron was fabricated into tools and weapons during the earliest times. Herodotus8 mentioned the use of iron in the construction of the pyramids. A number of iron specimens known to be of the period 2900-1450 B.C. have been recovered from the pyramid^.^ Diodorus as well as Plutarch described the ancient method of preparing swords and weapons of i r o n . ' q h e scheme frequently used by the Celtiberians in Aragon was to bury the piece of iron in such a way that it would corrode, "so that the weaker portions would be eaten away." This same method was used by the early Japanese. For a comprehensive historical survey of the metals and their properties, one can do no better than read C. Pliny's "Naturalis Historiae," volumes 33 and 34. Some of these data have been summarized by Weeks." Copper, which was known in the early times as meretrix metallorum, was most frequently called aes. Verdigris, or aerugo, was made by exposing copper to the fumes of vinegar. The well-known Dutch process for the manufacture of white lead was known and used by the early Romans. It is interesting to note that they were also familiar with the methods of adulteratine this material by additions of chalk, carbonates, etc. Pliny realized that the value of gold was due only to its permanence in resisting fire and the elements. He states, "In addition to this, gold steadily resists the corrosive action of salt and vinegar, things which obtain the mastery over all other ~ubstances."'~ Observations on silver were, "Silver admits also of being blackened with the yolk of a hard-boiled egg,"Ia and again, "Silver becomes tainted by the contact of mineral waters and of the salty exhalations from them."'4 In book 34, chapter 19, occurs a choice passage indicating that Pliny would have made an ideal member of the S.P.W., "Nature in conformity with her usual benevolence has limited the power of iron by inflicting upon it the punishment of rust; and has thus displayed her usual foresight in rendering nothing in existence more perishable than the substance which brings the greatest dangers upon perishable mortality." Probably the first reference to stress-strain corrosion is a description of the bridge constrncted by Alexander SWANK, "Iron in all ages." Swank, Philadelphia, 1884, p. 1. C ~ E N T EAND R ROBERTSON. J. Iron & Steel Inst., 121,
420 (1930). '0
l'
"
"
SWANK, 10c. cit., p. 16.
WEEKS,J. CHEM.EDUC..9, 4 (1932). PLINY, Book 33, Chap. 19. PLINY. Book 33. Cham 46. PLINY; Book 33, Chap. 46.
~~.
over the Euphrates. The bridge was held by iron chains, and it was noticed that the replaced links always corroded more readily than the old links.15 White lead mixed with tars and pitches was used to protect iron members from corrosion.'6 Copper articles were coated with tin by dipping the metal in molten tin. Articles prepared in this way were called incoctilia (inboiled). Pliny states, "When copper vessels are coated with stannurn they produce a less disagreeable flavor and the formation of verdigris is prevented." Vitruvius described chronic lead poisoning-the result of drinkmg water which had stood in lead pipes." Zosimus is credited with being the first to note that copper will precipitate upon iron when the latter is placed in a copper solution.'s Paracelsus in his "De Tinclura Physicorum" pronounced the deposition of copper on iron from solutions of copper sulfate to be a transmutation of the iron to copper. One can but imagine what excitement would have prevailed had some one snbstituted gold solutionsfor copper solutions. After Paracelsus, a considerable amount of information upon the action of various substances on metals was obtained in the Middle Ages. This plethora of data was no doubt due to the wide interest in alchemy, the prevailing science of the time, which directed attention to investigations of metals with the hope of a possible transmutation of the base metals to gold or silver. The prototype of the electromotive series of the elements was due to Bergman in the 18th century. In his "De praecipifutis metallicis" he states, "the metals precipitate one another after a certain order . . zinc prevails over iron, iron over lead, lead over tin, tin over copper, copper over silver. . . " In the 17th and 18th centuries peitung (white copper), frequently called packfong or packtong, was imported from China. These beautiful alloys owed their value to their comparative immnnity from tarnish. These metals are now known as the "German silvers." From 1700 on, the knowledge of corrosion and corrosion prevention increased rapidly. Climaxed by Lavoisier's work on the action of oxygen on ironlQ and other investigations, the modem age of chemistry came into existence. Today the man in the street knows the progress which has been made and is being made. Skyscrapers, such as the Empire State and Chrysler buildings in New York City, are covered with "stainless steels"; chromium-plated and "staiuless-steel" fittings are no longer novel for the automobile, kitchen, and lavatory; other articles are being plated with platinum metals and tungsten; scores of alloys of complex composition and treatment are constantly being presented on the market. " PLINY. Book 34. Chap. 43. " PLINY,Book 34, Chap. 54. " STILLMAN, "A story of early chemistry,'' D. Appletan & Co., New York City. 1927, p,. 23. IS MELLOR."Comprehensrve treatise of inorganic and theoretical chemistry," Longmans, Green & Co., New York City, 1923, vol. 3, p. 14. 'O LAVOISIER, Ann., 1, 260 (1789).
.
times called the electrochemical series of the elements, the electrode potential series, the electromotive series, the oxidation-reduction series, the activity series, etc. Table 1 lists some of these elements, the figures being taken from books by Latimer and Hildebrand and by Kremann and M i i l l e ~ . ~ ~ According to the convention of this table lithium is supposed to go into a oneTABLE 1 THE ELECTROCHEMICAL THEORY OF CORROSION Li , -2.g~ molal solution of its own ions, leaving a negative charge on the remaining metal A number of theories have been advanced on the NKn -2.92 -2.71 -2.70 of 2.96 volts. Copper, silver, and gold, cause and nature of corrosion, but the electrochemical C . although soluble in many media, do not theory has served to explain in the best manner all the : exhibit this solubility in one-molal solufactors involved. For an extended description of the Mn -1.00 tions of their ions; they all show a tenCr -0.76 other theories one can refer to any of the books on zn -0.46 dency to go out of solution, consequently corrosion. An excellent review of the electrochemical Fe Cd -0.40 -0.23 bringing their positive charges to their theory is given by Banaoft, one of the early champions ~i Sn -0.14 electrode. of the theory.*o Pb -0.12 Every one is familiar with the fact that H -0.00 The essence of this theory is that the solution of a ' an iron knife placed in a copper sulfate metal is the result of an electrochemical action at the As f 0 . 8 0 solution will become covered, more or less. surface of the metal. It is presumed that as an atom A= + i . a with a coat of copper. This experiment of the metal goes into solution it loses an electron shows that iron has a tendency to replace copper as (or electrons) and acquires a positive charge (or charges), forming what is called an ion. The remain- ions. The fact is that if the iron and copper could be ing, undissolved metal is then left with the excess elec- placed in a solution containing one-molal iron and tron (or electrons), giving it a negative charge. Now, one-molal copper the iron would have a tendency, this negative charge may be employed to neutralize a according to the table, to go into solution with a positively charged atom of another species present so potential of 0.46 volt and the copper to go out of the that the total action is that of one element going into solution with a potential of 0.34 volt. The total solution (losing its electronsbeing oxidized) and driving force of the action would then be the sum of the another element going out of solution (gaining its potentials, 0.80 volt. A galvanized iron pail behaves similarly. If the pail electroebeing reduced). The action is, therefore, is scratched so that iron and zinc are both exposed to one of chemical and electrical transformations. external action one might expect them both to go into THE ELECTROCHEMICAL SERIES OF THE ELEMENTS solution, but fortunately this is not so; only the zinc In order to avoid a philosophical argument on why a goes into solution, because it is easier for the zinc to go metal or element goes into solution, we shall merely say, into solution than the iron. A glance at the table then per scienter, that such a tendency exists to a greater or tells one which elements should go into solution easily lesser degree. A great deal has been written upon the and which with difficulty. Some of the elements, such solution pressures of the elements (tendency to go into as aluminum and chromium, exhibit non-tarnishing and solution), but the fundamental reason why one element non-rusting properties, so that they appear out of place; goes into solution faster than another is unexplainable, but these are properties, not of the elements aluminum and chromium but of their oxide films which cover the like chemical affinity. metal, and since these films are very resistant to most Since some of the elements go into solution more corroding mediums the metals act as if they belonged readily than the others, it seems reasonable that a table with copper and silver. The table does not misintercould be constructed showing the comparative solution pret; it measures the respective solution pressures of pressures of the various elements. Such a table can be the elements, not of their oxidefilms. constructed, and most chemists are familiar with it. If the reader will examine differentelectrode potential The standard for comparing the elements is hydrogen; tables he may lose all faith in electrochemists; some of in a solution one-molal with respect to the hydrogenthe tables give gold with a positive potential and some ion content, hydrogen is supposed to have a zero powith a negative potential and, conversely, the metals tential; that is, it has a tendency to go neither into above hydrogen with a negative potential or a positive solution nor out of solution. The other elements, being potential. The authors of one group state that lithium compared with hydrogen, fall into two classesthose and the metals above hydrogen are negative because, which have a tendency to go into a one-molal solution as those metals go into solution, the remaining metal of their ions, and those which have a tendency to go out must retain the negative charges; thus, they should he of such a solution. Since these tendencies are measured LATIMER AND HUDEBUND, "Reference book of inorganic in terms of volts and the tendencies are all different, chemistry," The Mamillan Co., New York City, 1929, p. 367; a table is verv easilv constructed. This table is some- OSTWALD-D~ucmn, "Handbuch der allgemeinen Chemie," On the other hand, the reader attempting to understand the many phenomena involved in corrosion is bewildered by the terms abounding in the technical and scientific journals-such terms as local action, polarization, inhibitors, stray currents, oxygen concentration, pH, electrode potentials, film formations, and other shibboleths. What do these terms all mean?
-
------
,
AND MWLER,''ElektTom~tori~che Krafte." vol.7, p. 2; KREMANN Akademische Verlagsgesellscbaft, Leipzig. 1930, p. 786.
negative. The other group argues that since the metals have a tendency to go into solution this spontaneity should be honored with a positive sign. To come to some sort of agreement the American Electrochemical Society went on record, after considerable thought, discussion, and debate, as adopting the metals above hydrogen as negative and those below as positive.z2 The same convention has been adopted by the majority of scientific people here or abroad, so that this nomenclature will be used in the present paper. The elements below hydrogen are sometimes called electropositive elements or noble elements, while those above hydrogen are called electronegative or base elements. Kahlenburg chooses to classify them as more zinc-like" and "more carbon-like." Convention dictates that nickel can be called more electropositive or more noble than iron or zinc, zinc more noble than magnesium, copper more electronegative than silver, etc. '6
It is sometimes possible for an electronegative element to plate out on an electropositive element, such as zinc upon copper. It is a well-known fact that copper will usually deposit upon zinc when zinc is placed in even extremely diluted solutions of copper. When cyanide solutions are employed, however, the activity of copper reduces itself to such a degree that the copper potential becomes negative and passes that of zinc. When this happens, and any copper metal is present, the zinc in solution will deposit upon the copper. The phenomenon of the plating out of one metal on another from solutions is sometimes called cementation by the electrochemists. This property is often realized in certain stages of the electro-refiningof metals such as nickel, copper, zinc, and lead. How intimately the electromotive series and the Nernst equation are connected with corrosion phenomena can be demonstrated by referring to Figure 1.
THE NERNST EQUATION
The electrode potential series gives us a fair indication as to the comparative corrodibility of the various elements and an exact relationship when the elements are compared in one-molal solution of their ions. When the solutions are not in the order of one-mold, or the concentrations are very diierent, then we cannot adhere too strictly to the table. It can be reasoned that as the number of ions of a solution decreases, the tendency for a metal to go into solution would increase; and if the ion concentration increases, then the tendency to go into solution should decrease. This reasoning can be proved experimentally as well as mathematially.^^ In 1888, Nernst developed a valuable equation to show the relationship between electrode potential and concentrati~n.~'Disregarding the development and derivation of his equation, which may be obtained from any standard text on physical chemistry, it may be stated as, E
=
Ea
Rt + nF - log C
(1)
where E is the potential a t some concentration, C, which is in terms of activity, Eo is the electrode potential as obtained from the electrode potential series, R is a constant, t the absolute temperature, n the valence of the element, and F the value of the Faraday (96,494). At 25'C. equation (1) resolves itself into E = Eo
+ +O
log C.
A zinc electrode is placed in a solution of zinc sulfate which is separated by a porous membrane from a copper electrode immersed in a molal solution of copper sulfate. According to equation (I), the zinc potential will be Zn = -0.76
If the concentration activity is less than one, the log term results in a negative value, thus making E less than, or rather, more negative than, EQ. If the concentration activity is more than one, then E will be more positive than EQ. It has been proved that the agreement between the calculated and the experimental values is quite close.
-
+ Og1nog
Czn++
and this resolves itself in a one-molal solution to Zn = -0.76 volt. In a like manner, the copper potential becomes Cu = +0.34 volt. When the two electrodes are connected with a wire, a current will flow with a potential equal to the algebraic difference of the single zinc and copper potentials, which is 1.10 volts. It makes no difference whether one wishes the final potential to be nepative or ~ositive:the im~ortantfact is that a difference of the two pot&tials is kserting itself.
If the zinc and copper concentrations are other than one mold, the following equation results:
E
=
-1.10
Co.++ - 0.059 -lag 2 cz,++
(3)
so that if the copper activity concentration is greater than that of the zinc, a greater potential than 1.10 volts will result. On the other hand, if a greater zinc than copper concentration exists, a smaller potential will result. It is a well-known fact that most corrosion is due to acid action, or simply a cell existing between hydrogen and a metal. Thus, if we had substituted an acid for the copper solution in the previous illustration and a hydrogen electrode for the copper electrode, our potential would be: E = (-0.76
+ 0.059log &++)
- (0.00
+
log CE+
or: E
=
-0.76
2Cw - 0.059 log 2 cz,++
(4)
I t can be readily seen, then, that the concentration of the acid has a great deal to do with the corrosion rate. With metals like nickel and tin, which are very close to hydrogen in the potential series, and with very weak acids, no appreciable potential should be present, and consequently no corrosion results, cczteris paribus. HYDROGEN POLARIZATION AND OVERVOLTAGE
It may be difficult to conceive of a hydrogen electrode, but such an electrode is quite possible. By definition, the hydrogen potential is zero in a solution with an activity of one-molal hydrogen ions. Hydrogen goes into solution, or out of solution, in two distinct steps 2H+
+2r=2H"GH9"
(5)
which are the ionic to the atomic to the molecular, when it is going out of solution, and the molecular to the atomic to the ionic when i t is going into solution. The change from the atomic to the molecular state is zero when the process takes place a t the surface of the platinized platinum electrode; a t the surface of other metals, however, there is generally an appreciable energy change attending this transfer. As a rule, molecular hydrogen will deposit out as bubbles a t the surface of metals only a t the expense of extra energy which can be measured in terms of volts. Platinized platinum, palladium, and graphite seem to give the only surfaces which permit the evolution of hydrogen without expense of energy. All other metallic surfaces have a tendency to resist the change of atomic hydrogen to molecular hydrogen. This resistance is a form of hydrogen polarizetion. The nature of the metal surface, the concentration and temperature of the solution, and the current density, are factors
which influence the degree of hydrogen polarization; but since these are secondary factors, they may be disregarded for the time being. The elements can be arranged in the order of their tendency to resist the evolution of hydrogen a t their surfaces by comparison with the evolution of hydrogen a t the surface of platinized platinum. This value measured in terms of volts is called the hydrogen overvoltage of that element. Thus, hydrogen polarization and hydrogen overvoltage are very closely related; the former represents the phenomenon and the latter has, within limits, a definite numerical value. Table 2 gives a list, taken from N e ~ b e r y of , ~these ~ values for some of the more common elements. One can roughly say that the soft metals, such as zinc, lead, and mercury, have a Zn = 0.70 high hydrogen overvoltage, while the HZ - 0.61 M= = 0.58 hard metals, such as nickel, iron, and TI -0.55 cd = platinum, have a low hydrogen overpb 0.45 voltage. This hydrogen polarization Sn = 0.44 cr = 0.41 manifests itself in corrosion reactions '0.39 in several ways. Sb = 0.38 A" = 0.38 Every one familiar with the preparation C = 0.34 of hydrogen from zinc and acids realizes = Ag 0.33 that any pure metal stubbornly refuses Ni = 0.29 to go into solution and displace hydrogen. = o.z8 = 0.24 The reason becomes quite apparent when Co = 0.23 pt = 0.20 one studies the potentials involved. Ac0.18 cording to equation (4) the potential will Pd = 0.00 R-~t 0.00 be in the order of 0.76 volt, while according to Table 2 the hydrogen overvoltage will be in the order of 0.70 volt, leaving a total of 0.06 volt to drive the reaction forward. One is safe in assuming that when the reaction is starting, the zinc concentration will be so low and the hydrogen concentration so high that a sufficient electromotive force will be generated to drive the reaction toward hydrogen evolution; but after the zinc surface has been attacked by the acid, the rate of diffusion of the zinc away from the active surface is so slow that its concentration asserts itself and the reaction immediately slows down. Before hydrogen can be evolved, this equation must hold:
-
-
Where E, is the potential of the metal, E, the potential of the hydrogen, and Eo the overvoltage of the hydrogen on the metal. What is frequently done to promote this reaction is to add a few drops of copper sulfate solution. The copper deposits out on the zinc, as previously explained, and the evolution of hydrogen begins immediately because the overvoltage of hydrogen on copper is of the order of 0.33 volt, and then, according to equation (6)
the reaction can proceed with a potential of 0.43 volt. Eventually, the zinc concentration of the solution will become so high and the acid concentration so low that
7. Ckm.
SOC.,
109,1051 (1916)
even this reaction will stop. Addition of traces of metals like nickel and platinum, with their lower overvoltages, should accelerate the reaction even more. That this is true can be ascertained from the experimental work of Centnerszwer and S a ~ h s . ~ ~
Solut~onrate increas~nq
/
olut~onrote af
THETHEORYOFLOCALACTION
The explanation of the galvanic action of metals during the corrosion of one of the elements was first made by De La Rive2' and has since been called the theory of local elements, the theory of galuanic elements, or simply, local action. Graphically the theory can be explained by Figure 2. been uncovered, reducing the active atea of the dissolving metal. . Section (C) shows the surface of the metal nearly covered with the impurity. At this point, the reaction has slowed down, and it is even possible that the impurities would be attacked by the acid. The period from the time the metal is placed in contact with the corroding solution until the solution velocity is a t its maximum is called the period of induction. The significance of this period has frequently been overlooked in the corrosion researches in this country. Its application will be discussed under an appropriate heading later. AurerP8 and Palmaer2*have contributed very valuable information on the theory of local action and its application. Not only were they interested in applying the Nernst equation but they also developed an equation involving the possible electrical resistance of the local action circuit. The derivation of their formula is very instructive and simple. Their problem was to find the velocity of the reaction between a metal and an acid in terms of the number of cubic centimeters of hydrogen evolved in unit time; then if 4 equals the number of cubic centimeters evolved in one ampere minute, i will equal the total intensity of the local action currents:
A piece of metal, say zinc, contains some impurity, such as iron. The metal is placed in an acid solution and, since there is a greater tendency for the hydrogen to replace the zinc than to replace the iron, only the zinc will he attacked. The first stage of the action is represented by ( A ); the hydrogen ions are replacing the zinc and discharging a t the surface of both the zinc and the iron as in (B). It will be noted that the action is confined to the interface of the liquid, zinc, and iron. Since it is easier for the hydrogen atoms to evolve a t the surface of the iron than of the zinc (the overvoltage being lower), hydrogen evolution will commence on the iron. This action continues with the formation of a cavity a t the zinc-iron interface until finally all the zinc is corroded from the iron surface and the iron particle is dislodged, leaving a pit. If the piece of iron is too dV/dT = 4i. heavy to be carried out of the cavity by the liquid or gas currents, it remains and sinks deeper into the zinc as And since E = i R where R is the resistance of the local the zinc is being dissolved. The final result may be elements and E the potential of the local elements, as shown in (C). If the iron particle had been washed away, then the action would cease unless other iron particles were present. What actually happens is that more impurities are uncovered as the corrosion One can denote the specific electrical conductivity of action continues, so that more local action is encouraged the solution by L and the resistance capacity of the and the corrosion rate speeds up. This increase in the local elements by C, so that R = C/L; then, reaction velocity soon reaches a maximum and then drops down again because the active zinc surface has been gradually replaced by the non-dissolving iron (or other) impurities. This may be represented as in If we take the value for 4 under ideal conditions and Figure 3. Figure 3 ( A ) represents a side view of a 760 mm. and O°C., the equation gives, metal which contains impurities being attacked by a corroding solution. Section (B) shows the metal after i t has corroded and three impurity particles have
-
" CENTNERSZWERAND SACHS. Z.p h y ~ kChem., . 87,692 (1914). " Ann. chim. phys., 45,425 (1830).
If, for example, iron is being corroded with an acid, the following will hold: E = EP* - Ea - Eo E = 0.46 - 0.24 0.22volt
-
Then.
0.059 6.95L(0.22) - 7 (log Cp.- 2 log Cd dV/dT =
C
All the factors can he obtained in advance except C, which can be established as a constant after several runs on the same metal. Palmaer in his three-volume series on corrosionao gives a splendid account of the theory and application of local action. It should he appreciated that most of his results have checked his theory time after time. When we discuss the factors affecting corrosion i t will become more apparent just how valuable equation (7) is in explaining many of the factors. IMPORTANCE OF OXYGEN
So far only those corrosion reactions involving the evolution of hydrogen have been discussed. With most corrosion reactions the hydrogen-ion concentration is so low that i t is impossible for a great enough potential to be developed so that hydrogen is evolved. Take, for instance, the corrosion of iron water pipes and outdoor iron and steel work which are subject to the action of water. Disregarding the effects of dissolved gases and salts, the hydrogen-ion concentration of water will be around lo-' gram ions per liter. This concentration is extremely low, and considering the overvoltage one can see why hydrogen evolution does not proceed. If the overvoltage is very low, then it is possible that hydrogen .evolution will take place. The author has taken palladium and platinum-plated base metal specimens, placed them in distilled water, and observed hydrogen evolution a t the scratches. Fortunately, most metals do not contain platinum and palladium as impurities, s o that this evolution of hydrogen does not take place. If hydrogen were not removed from the surface, ,corrosion action would cease or would be so low as to be insignificant. Oxygen is always present in solution, however, and has the ability to react with the atomic hydrogen, forming water. This, then, is a means whereby the discharged hydrogen is constantly removed, resulting in a slow hut steady rate of corrosion, the rate being dependent, of course, upon the diffusion of the dissolved oxygen to the active metal surface. Oxygen in this rBle is called a depolarizer. Equation (7) then be~omes,~'
where dV/dT is the reaction velocity in terms of the amount of metal dissolved in unit time, the amount of metal dissolved per minute, Em the potential of the metal according to Table 1, EO the overvoltage, and Ed the potential of the depolarizer; the other terms are the same as before. The absence of oxygen, then, in all cases where hydrogen does not evolve, results in negative corrosion reactions. Too much oxygen, on the other hand, may promote the formation of insoluble oxide films which inhibit attack. These slightly soluble oxide films are called passive films and the phenomenon is known as fiassitity. It is evident, then, that too much or too little oxygen inhibits corrosion while a definite amount promotes it. This brings one to an enormous field of study-that of the influence of oxygen concentration. With the metals below hydrogen in the electromotive series, such as copper, the oxygen forms an oxide film which is soluble in most solutions. In this rble, oxygen acts in a primary way, attacking first the metal, and then the resulting oxide, which dissolves in the acids or salts in solution of the corroding medium. In this section only those corrosion phenomena have been studied which exist under so-called ideal conditions. Any one familiar with the actual facts realizes that there are a very great number of factors which (individually and collectively) influence the corrosion of metalssuch as kind of solution, physical state of the metal, temperature, concentration, oxygen concentration, diffusion, light, colloids, etc. These factors are t o be discussed in the following paper. (Part II of this series m'I1 afifieer in the April issue.)
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STUDENT ATTENDANCE AT THE WASHINGTON MEETING OF THE AMERICAN CHEMICAL SOCIETY
The American Chemical Society makes a special concession to regular students, of either colleges or high schools, by allowing them to register and attend its meetings a t the minimum rate of $3.00. There will doubtless be many students who will want to attend the meeting of the Society taking place in Washington, D. C., during the week of March 27. The Wasbington Y. M. C. A. has agreed to furnish accommodations for such students a t the minimum rate of 81.00 per day in double rooms or $1.50 in single. The Division of Chemical Education is always anxious to help students log cm to attend the Society's meetings and through its seae~ W T= 6 ( - E , - EO E~ - 0.059 - - h c H ) ) tary (N. W. Rakestraw, Brown University, Providence, (8) R. I.) will be glad to correspond with any who wish a PALMAEX, "The corrosion of metals," Ingeniorsvetenskeps- further information concerning details. In certain akademiens Handlinpar. No. 93 and 108. Svenska Bokhandelscases in the past it has been ~ossibleto arrange transcentralen, 1929. PAL.MAER, lo
