Corrosion: A Waste of Energy by J. CHEM. EDUC. Staff
Broadly speaking, corrosion is defined as the deterioration of a substance by environmental constituents. Thus, although technically the attack on a limestone building facade or marble statuary by "acid rains3'as well as the rusting of automobile bodies are examples of corrosion, we limit our discussion here to the attack on metals by the environment. Common, naturally occurring atmospheric constituents which contribute to the formation of a corrosive environment include oxygen, carbon dioxide, and water vapor as well as the many natural waters which contain a variety of dissolved solutes. If the usual industrial pollutants, e.g., sulfur oxides and nitrogen oxides, were not present in our environment, we should still be faced with the oroblems associated with metal corrosion. Indirectly the corrosion of metals represents a serious energy drain in an industrial society. Energy is required to produce metals from ores, and the loss of any significant amount of these metals that return to the ore-like state through the corrosion process means that additional metal must he produced a t the exoense of more enerm. For exam~le,it has been estimated thatabout 25% of the world's production of steel is lost bv corrosion. In this countw alone, which produces about 26% of the world's steel, theloss by corrosi& corresoonds to about 32 million tons of steel annually-a not inconsiderable amount in terms of extra ene& requirements.
Thermodynamic Aspects of Corrosion Corrosion processes (eqn. (1)) involve the spontaneous conversion of metals or alloys into their compounds-carhonates, sulfides, sulfates, XM(C) + YZ- M,Z,(C)
(1)
hydroxides, oxides, or hydrous oxides. We know from thermodynamics that spontaneous reactions must have a negative free energy change (AG), which, in general, can be calculated for any reaction to be Recall that the symbol AGrn means the free energy of a substance in its standard state, which is normally its state at 25°C and 1 atm pressure. Thefree energy of any element in its standard state is taken to be zero. To illustrate the importance of the stability of the corrosion product in driving reactions like (1) to completion, let us consider the reaction of typical metals with oxygen (eqn. (3)). Since the free energy of any element XM(C) + Y1202
-
M,O,(C)
(3)
in its natural state is zero.. e m s the tendencv . . (2) . ~ r e d i c t that for the process in eqn. (3) to occur is driven by the free energ; of the oxide that can be formed. A consideration of the free ener.a!e of the oxidm of tht' common structural metals t'l'ahlc 1) indicates that a11 would go svontaneou.ily to the oxide in an atmosphere ot'uxygcn. Th& ihermodynnhics indicates that the mc.tallir state is nut ~~rcferred energetiuilly in our world, and hard won metals will attempt to revert 10 the or(:-like state. Similar thermodynamic arguments indirate that the f n r m a t ~ mof all the pnldurt.; formed in corrosion processes (eqn. ( I )) is energetically more favored than the metallic state. Although thermodynimics tells us whether a reaction will (rcur, we cannot deduce the rate at which reactions occur fmm these arguments; fortunately, kinetic factors intervene in corrosion processes, which blunt the inevitable results of the thermodynamic considerations.
Table 1. Free Energies of Oxldes of Common Structural Metals Oxide
AG,~,kcallmole
F e A A1203 Cr& MgO CuO NiO ZnO Sn01
-177.4 -378.2 -252.9 -136.1 -31.0 -50.6 -76.1 -124.2
Table 2. Standard Reduction Potentials Reaction'
P , volts
Au3+(aq) 3eAu(C) 0&) 4H+(aq) 4e2HpMO Agt(aq) eAg(C) 2Htq1) 4e40H7aq) 0&) Cu2+(aq) 2eCu(C) Z H + ( ~ 28H&) Sn2+(aq) 2eSn(C) 2 6 Ni(C) NiZ'(aq) Fe2+(aq) 2eFe(C) C++(aq) 3.eCr(C) ZnZt(aq) 2eZn(C) APt(aq) 3e' AI(C) Mg2+(aq) 2 8 Mg(C) Natlaq) eNa(C)
1.50 1.23 0.80 0.40 0.34 000 -0.14 -0.24 -0.47 -0.73 -0.76 . -1.68 -2.36 -2.71
+ +
+
+
+
+ + + + + + + +
+
- - -+ +
-----
he state of the species ie indioaled by (Cl, uyrtalline solid: (I), liquid: (aq).aqusted
ion; (gl, gas.
Electrochemical Aspects of Corrosion There is considerable evidence that corrosion processes are fundamentally electrochemical in nature. Note, for example, that the corrosion process described by eqn. (2) is one involving oxidation and reduction. Accordingly, it should be possible to discuss corrosion processes in conventional electrochemical terms. Recall that all redox processes are comprized of two half-reactions, each of which is associated with a characteristic potential; the potential of the overall redox process is the algebraic sum of the appropriate half-cells. The familiar relationship between the free energy of the redox process and its potential (eqn. (4)) indicates that a positive
AGO = -nFE"
(4)
cell potential corresponds to a negative free energy, and thus, a spontaneous reaction. In eqn. (4), n represents the number of moles of electrons transferred, F the Faraday constant equivalent to one mole of electrons (96,500 coulombs), and En the standard cell potential, which is determined for aqueous solutions of unit activity (approximately 1M )a t 2 5 T and 1 atm pressure. From the point of view of electrochemistry, standard potentials (e.g., Table 2) become a source of information concerning corrosion processes. For example, the half-cell readions in Table 2 predict correctly that iron would be oxidized in preference to copper because the reaction E0= +0.47V F d C ) Fe2+(aa)+ 2e-
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Volume 56, Number 10, October 1979 / 673
potential is positive, which leads to a negative free energy. The elements of this cell are found in many plumbing systems where a steel pipe is attached to a copper pipe. Except for the difference in the ionic activity between the standard state and ordinary tap water, the thermodynamic arguments lead to the svontaneous dissolution of the iron nine. . Usine~.a ealvanized pipe vinc coated steel) does nut inipn~vethe situatiun l~ecause thr uutrntial of the zinc half-cell ('l'ahle 2) issut'ficientlv.high to cause it also to dissolve in contract withcopper. Even if two metals are not joined, possibilities for corrosion exist if the metal is in contact with an aqueous solution-a common occurrence in many parts of our environment. If a metal is to be corroded (oxidized), there must be a corresponding reduction process. Two such processes are possible in aqueous solution (Table 2)
.
-
Air 02
F M"+
O2*OH-
4
aqueous electrolyte
-
--
EO = O.OOV (5) 2H+(aq)+ 2e- Hz(g) 0Ag) + 2H20(1) + 4e40H-(aq) Eo = 0.40V (6) In neutral solutions, the reduction of H+(aq) (eqn. (5)) is negligible. "Acid rains" are reasonably common occurrences in urban areas and, accordingly, in such environments metals can corrode readily. The overall process would consist of
2H+(aq)+ 2eM(C) 2nHt(aq) + 2M(C)
--
En = 0.00 (7) Hz(g) Mn+ + neEM' (8) 2nHz(g) + 2M"+ En= Emo (9) Any metal with a positive oxidation potential, so that EO for the reaction is negative, would corrode in an acidic environment. The arenment. as usual.. ienores the fact that ionic ac" tivity in the natural waters is not the standard value. According to this argument all metals that are more active than hydrogen (having negative reduction potentials, Table 2) would corrode in such an environment. Ins~ectionof Table 2 indicates that many of the common metals of construction are included in this class. Ewn in the nhsrncr ofexcrss H + ( a q ~the , reduction of 0 2 leqn. ( 6 ) )provides the potenlial for the other half-cell in metal corrosionprocesses. ~ h n sa, piece of metal (e.g., iron) in contact with an aqueous electrolyte saturated with oxygen undergoes corrosion via an electrolytic process (Fig. 1).The oxidation of the metal, which is an anodic process, leads to formation of Mn+(aq). The electrons which are liberated during the metal oxidation find their way to appropriate surface sites (cathode) where reduction of oxveen can occur (eqn. (6)). ~ o t i c that e the solution around tde-cathodic site becomes basic. If the metal forms an insoluble hvdrous oxide arr hydn~xide,it is possil~lefur the product to precipitate at a ooint in solutiun which is well seoarated f n m either the raihodic or anodic site as the ions diffuse away. Of course, additional chemistry can occur in solution if the metal ion which is initially formed can be further oxidized, as is the case of iron. The primary ion formed in the corrosion of iron is Fe2+(aq),which undergoes further oxidation by dissolved oxygen to give Fe3+(aq);it is the latter ion which precipitates 3s the hydrous oxide, Fe20a.xHzO(rust), when iron corrodes.
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Inhibition of Corrosion Processes If the thermodynamic arguments lead to the conclusion that, e.g., metals will be eventually ~:onvertedto these oxides. tvhg is it that we can still use very active metals like magneiium and aluminum for xtrurtural components? \Vhy doesn't I modern airvlane. which contains a liree number of alnminum strnctnrh elements, quickly become a pile of aluminum 3xide as the data in Tahle 1suggests? The answer lies in the nature of Alz03, which is isomorphic with metallic aluminum; b a t is, there is a crystalline continuity between the first layer ,f aluminum atoms in the metal and the first layer of A1203 ~ h i c hforms upon it. In effect. the A1.0~ forms an imoervions 'skin" on the-surface of themetal is soon as i t cokes into :ontact with air. The diffusion of oxygen molecules t o the 174 / Journal of Chemical Education
M-Mn+ he processes involved in
+ na-
O2 t 2 H 2 0 + 4s--40H-
metal corrosion in aqueous electrolyte solutions.
surface of the metal is greatly impeded by the A1203 SO that the cathodic process does not occur. Since Al& is airtually insoluble in water, placing metallic aluminnm in contact with water (ex., an aluminnm sauce Dan) does not increase corrosion. However, if the aqueous soiution contains chloride ions, as in the case of seawater or brines, corrosion of aluminum occurs more rapidly. Chloride ions are believed to enter the structure of AlpO3, creating defects which lead to the rupture of the previously impervious A1203 skin; oxygen (or water molecules) can attack the aluminum substrate a t these defects, which leads to localized corrosion (nittine). -, The formation of an outer passive coat, fortunately, occurs with several structurally important metals. Magnesium forms MgO, zinc forms a basic carbonate [ Z n C 0 ~ n Z n ( 0 H )in ~ ]air rather than an oxide. as does comer and .. [CuCO&n(OH)?l. . chromium forms ~ r i 0 3 .The passive natuie of ~ r z 0 ;has several advantaees: i t takes a hiah and thus is useful . volish . from a decorativ~poinr of view, and alloys containi~~g chn,mium (stainless steelj ;are protected because the CrSL is extremely durable and readily repairs itself when the film is broken if oxygen is available. The hydrated iron oxides are soft and generally non-adherent; they encourage further corrosion by occluding moisture and electrolytes, which is essential for the electrochemical cell that is active in the corrosion of iron. Although the hydrated iron oxides themselves are not passivating, it is possible to artificially produce a passive skin on iron-containing items bv oxidizine them with concentrated HNO?. aaueous I(-2Cr207,ormolten KN03. This procedure forms a iightly adherent skin of iron oxide (Fen031 on the iron surface; thick blue-colored layers of this oxide ("gun blue") can be estahlished with repeated treatment by molten KN03. Iron treated with dilute solutions of strong oxidizing agents becomes more noble by nearly 1volt (from -0.47 V in Tahle 1to +0.60v for iron with an Fez03 coat); such solutions are often used as corrosion inhibitors in steel systems containing circulating cooling waters, as in internal combustion engines. An understanding of the electrochemical nature of corrosion has led to several simple techniques for inhibition of the key processes. Thus, since the cathode reaction in the corrosion process [Fig. I I often cannot Ir: eliminated (eg., cstaldishment of complet~~ly anhydrousor anaenhir conditions) fur pmvtsal reasons, a sacrificial anode is provided to protect themetal of interest. For example, zinc is a more active metal than iron and is used to coat steel items (ealvanized steel). . . which allows the zinc to corrode (and perhaps form a passive coat) in preference to the iron. A piece of magnesium attached by an electrical conductor to a buried iron tank serves the same purpose. A ship's steel hull can be protected in the same way by using massive pieces of magnesium or zinc; the latter substances oxidize in preference to the iron in the hull. It should be apparent that a metal object can be protected from corrosion if it can be snpplied with excess electrons (to reduce either oxygen or Ht(aq) (eqns. (6) and (7)); the sacrificial anode serves this purpose. A similar effect can be achieved by impressing a low-voltage direct current on the object to he protected. Such currents can be supplied by small generators, often powered by windmills in isolated areas.
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