California Association of Chemistry Teachers W. H. Slabaugh
Corrosion
Oregon State University Corvollis, 97331
The topic of corrosion is usually considered to he too mundane and dull to include in an academic chemistry course. and vet it extends several basic concents of electrochemistry, crystal structure, and kinetics into practical chemistry with which the student can relate. Furthermore, there are several experiments and demonstrations which, if not s~ectacular. are dramatic illustrations of corrosion and hear out thefundamental chemistry of metals. Standard Electrochemical Potentials
In connection with electrochemistry and redox chemistry, the concept of the standard electrode potential is a common starting point. We frequently rely on a tahle of such potentials to stress the tendencies of the various species of metals, nonmetals, and their ions to enter into chemical reactions, even though it is often difficult to reproduce because these potentials are based on thermodynamic data rather than direct measurements in aqueous media. With this in mind, a list of standard electrode potentials is given in the table which includes several peculiar entries. For example, the typical Fe-Fe(II) potential of +0.44 V is accompanied with a similar "standard potential" for iron in the passive state. In the passive state, iron seems to be coated with a thin oxide layer which can be simply represented as Fe.O.Oz. That is, by inserting a piece of iron into concentrated HN03 there seems to be very little reaction because a passive oxide layer is formed. Uoon removine the iron from the acid. washine with water, and reinserting the passive iron into a solution of CuSOa. the passivitv of the iron is obvious because no reductiodof ~"(11)o c k s . Upon striking the surface of the passive iron with a sharp object (a file or a glass rod), the passive film is punctured and rapid reduction of the Cu(n) spreads to the entire surface of the iron. It is generally agreed that passive iron is coated with both an oxide layer and an adsorbed layer or layers of oxygen molecules, hence, the notation Fe.O.02 for passive iron. For more permanent protection of iron, a thicker laver of " a n blue" is needed. and this can be achieved bv immersing the iron in molten KN08 or similar strong oxidant. Such thicker oassive coatines are cornnosed of adherent, hard layen of oxides such a s ~ e ~ 0 ~ . Aluminum, according to the academic viewpoint, should never be used to build airplanes because they would dissolve in the first rainstorm. Fortunately, pure aluminum with a standard electmootential of +1.66 V. immediately forms an oxide layer when exposed to air and this laver of AlrOz. nronerlv - -. or more . . - denoted AMOH). .. .. because it is somewhat similar to a gibhsite layer, forms a c o m.~ a c.t .adherent, and hard nrotective surface coatine on the metal. ~onseq"ently, the-standard e ~ e c t r o ~ o t e n tof i~l aluminum drops sharply to nearly -0.60 V where it appears to be neaily a typical noble metal,
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Presented to the California Association of Chemistry Teaehen at Asilomar on August 24, 1973. 218
/ Journal ot Chemical Education
Other metals in common use which develop more passive states when they interact with certain environments include silver which forms a nrotective laver of AeCl and lead, in moderately concentrated sulfuric acid, which forms a nrotective laver of P ~ S O LIron . becomes about 0.5 V more noble when ;mmersed in 0.002 M K2Crz07, an indication of the effectiveness of such substances as corrosion inhibitors in cooling systems. When we move to alloys, such as the stainless steels, passivity again occurs merely in the presence of oxygen. That is, a typical Fe-Ni-Cr alloy exhibits a noble metal electropotential of about -0.60 V (see the table). A stainless steel ordinarily contains a minimum of 12% chromium which satisfies the formation of an intermetallic compound Fe&. Chromium, with five vacancies in its 3d subshell, supposedly ties up one of iron's 3d electrons, thus converting the iron atoms in stainless steel to the chromium electronic structure. Pure iron first dissolves to Fe(II), but iron from a stainless steel first enters solution as Fe(II1). Furthermore chromium, which is nearly as electropositive as zinc, forms a protective oxide layer like aluminum; hence stainless steels automatically become passive in the presence of oxygen. A sidelight on stainless steel's ability to vary between an active.. hiehlv electronositive metal and a nohle metal can he dramatically illu&rated. By placing a rubber hand around the neck of a stainless steel coffee pot (use a discarded pot or simply a flat piece of stainiess steel), the metal surface under the rubber hand will he oxygen-deficient compared to the larger, open areas. After two weeks of immersion in a dilute salt (NaCI) solution, the covered area (oxygen deficient) will have corroded away. As a rule, stainless steels should always be used in an oxygen environment and uniformly exposed in order to avoid the formation of active areas which lead to pitting and loss of the metal tosolution. The much slower corrosion rates of the noble metals such as copper and silver, is a reflection of their lower electrochemical potentials. Protective layers of corrosion product on these metals are, however, more elegant if not more comnlex than their oxides. In normal atmosnheric exposure, copper forms a basic copper carbonate, sometimes called patina and lends an air of distinction to older
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Standard Electrode Potentials; EMF Series
Electrode Reaction K%Kt+eMg 5 Mg" 2eAl e Als+ 3e-
+ + Zn % ZnZ++ 2,Fe t;j Feat + 2eNi % Ni" + 2eHZQ 2H+ + 2e-
+
Cu % CuZ+ 2eFe.0.02.(passiye)
. .
Fe-Cr-NI (passwe) A1.A120a Agt;jAg++eAu % AuJf 3e-
+
E" (V)
Fe+-~ef*+~e-
2~'+2i+H, Qt2Hz0+4e'+4OH~
Figure 1. The mechanism of iron corrosion.
buildings with copper roofs and flashings because of the pleasing green color of the patina. Silver develops a layer of AgzS, sometimes called tarnish, which in thin layers adds a note of warmth to an already warm, metallic appearance. Corrosion of Iron
The mechanism of typical iron corrosion has been generally conceded to he electrochemical in nature: Anode and cathode areas are involved: electrons flow from anode to cathode, oxidation of iron to Fe(I1) occurs at the anode, and several reduction reactions occur at the cathode. See Figure 1. At anode areas of iron, the iron is electrochemically oxidized to Fe(I1). In an oxygen environment, the Fe(II) is quickly oxidized to Fe(I1I) which is subsequently changed to Fe(0H)s and finally to a hydrated ferric oxide. The degree of hydration of the oxide influences the color of iron rust which may vary from black to yellow to red-brown. The intermediate Fe(II1) ion is, incidentally, a promoter of corrosion and the iron oxides are soft, non-adherent deposits that encourage further corrosion of the metal by occluding moisture and electrolyte, all of which are essential to the electrochemical cell that operates in the corrosion of iron. Corrosion of Aluminum
Lest we get the idea that aluminum does not ordinarily corrode in common usage, we are reminded that even the airlines spend much effort to overcome and control corrosion of their airplanes. One of the principal problems with aluminum is the presence of the chloride ion. Aluminum promptly forms an oxide coating which is uniquely isomorphic to the base metal. In Figure 2, one layer of aluminum atoms has been oxidized to AI(II1) and the size of the oxveen .- atoms in the oxide laver accounts for a favorable spatial arrangement so that there is essentially a perfect fit between the oxide laver and the base metal. The crystalline continuity between aluminum and its oxide is, however, seriously disrupted by the presence of an occasional chloride ion, as noted in Figure 2. Whenever such a defect occurs, the oxide coating is ~ p t u r e dand the rupture develops into an anodic area where further oxidation of the base metal occurs. Consequently, it is believed that the presence of the chloride ion is largely responsible for the corrosion problems with aluminum. A simple experiment that demonstrates the speed of aluminum corrosion when the oxide layer is disrupted can be easily performed. Place a drop of H g z ( N 0 h solution on an aluminum surface; rub the drop into the metal with a piece of sahdpaper; blot up the excess solution; and within 5 min a deposit of aluminum oxide up to 2- or 3-mm thick will develop a t that region where the oxide layer was destroyed and the mercury atoms, reduced from Hg(1) by the exposed aluminum, now prevent the formation of an adherent, compact aluminum oxide. Applications of Electrochemical Concepts
Figwe 2. Cross-section of crystal struct~re01 aluminum and its surface layer of aluminum oxide.
by imposing restrictions, such as polarization of anodic and/or cathodic areas or by superposing a counter voltage on the corrosion cell. Again, the rate of corrosion may he increased by supplying conditions favorable to the electrochemical mechanism. It is customarv to illustrate an electrochemical cell with the anode on tbk left, the cathode on the right, with electron flow in the external circuit from left to right. Thus, we freauentlv see in textssuch illustrations of the ZnICu galvanic cell" along with a salt bridge connecting the 'two half cells. Essentially the same kind of galvanic cell is often set up in many plumbing systems, whether intentional or not (Fig. 3). Here, a galvanized pipe (zinc-coated steel) is attached to a copper pipe and, except for the difference in ionic activity between 1M electrolyte and ordinary tap water, the resulting galvanic cell results in the spontaneous dissolving of the zinc pipe-a boon to the plumber. The flow oftapwater freq&tly depolarizes the anodic and cathodic areas by removing the corrosion products and supplying new electrolyte for the continuing cell reactions. In a more favorable application of electrochemical principles, we can prevent or inhibit the corrosion of an object by making the object cathodic to its environment. Such is the case when we "cathodize" a buried iron pipe or tank (Fig. 4). By selecting a more electropositive metal than iron to serve as an anode (Mg or Zn is commonly used) the iron is made cathodic and remains practically unaffected. The size of the anode should he such as to provide
E';
,, * . ,.
,o
a
.' Mq a,
;.:.LJ.:; ...._ \
If the mechanism of iron corrosion is tmly electrochemical. then we should he able to alter the rate of corrosion
1.10 vo1tr
Figure 3. A zinc-copper galvanic cell in a plumbing system.
1
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Anode
Iron
oil tank ..
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.
'
Cathode
Figure 4. The calhodic protection of e buried iron tank.
Volume 51. Number 4. April 1974 / 219
T, = 75'C then E = 0.004 v
+mjet im ib./i" 30
tt1.n
5.ulmr
EL :890 volts 1, = 5. to-'amp 00019 g r a m ~ l / h c u r
Figure 6. Streaming potential in aflowing water system. Figure 5. Galvanic potential in a water heater as the result of hot and Cold regions.
about 2 mA current for a typical buried oil tank. Soil resistivity is a controlling factor and should he taken into account. The magnesium anode is buried a t least 3 m from the tank and in a bed of sand which speeds up the diffusion of anodic products into surrounding groundwater. T h e two electrodes are connected with a suitable electrical conductor such as a copper wire. Alternately, a n inert graphite anode is commonly used along with a n externally imposed emf from a d c power source. Another practical application of electrochemistry exists in the home water heater. The Nemst equation is commonly applied to problems of variable electrolyte concentrations; whereas this equation also provides for temperature variations. In addition to the water heater being a "hot spot" in a plumbing system, the tank itself is subject to variations in electrochemical potential a s shown in Figure 5. In this application of t h e Nernst equation it is assumed t h a t the ionic concentration, the Q factor, is constant. One more application of electrochemistry in corrosion occurs in a n irrigation system. Whenever a liquid flows through a tube, a n electrical potential is generated hetween the two ends of the tube. This is called streaming potential and is the result of the continuous shearing of the diffuse double layer a t the liquid-solid interface hetween the flowing liquid and the tube. The tube may he either metallic or nonmetallic, hut if it is metallic, then corrosion is potentially likely t o occur. In Figure 6, a typical sorinkler is illustrated. delivering water a t rates indiE,, of 890 V and cated. A continuous streaming current. I,, of nearly 1 PA will account for the solution of the alum&um pipe-which, depending on the ionic nature of the irrigation water, explains why a farmer occasionally finds his irrigation pipes have disappeared. Streaming potential, incidentally, is dramatically illustrated in the modern jet aircraft engine where the fuel (kerosene) is forced into the combustion chamber under
220 / Journal of Chemical Education
very high pressures and rates of flow. S o high is the streaming potential t h a t a mntinuous spark is formed hetween the fuel jet and other engine parts; all of which takes place in the combustion zones of the engine. Some Experiments
In addition to several corrosion-related experiments suggested shave, two more are described here. Several factors can he explored which commonly influence the rate of cornsion. Perform these tests an a piece of cold-rolled steel sueh as fender stock which can be obtained at a local body and fender shop. Remove all grease, oil, and paint from the steel; then add a drop of dilute solutions of each of these materials to various areas of the steel panel: HCI, NaOH, H202, FeC13, FeSOn, KzCr20r, NaNOz, NaCI, and any others the student might wish to investigate in this way. Observe the time at which evidence of a reaction occurs (evolution of gas or formation of a rust deposit). Some of these materials encourage corrosion; others inhibit corrosion. Then use combinations af these materials sueh as an active corrosion stimulator along with a corrosion inhibitor. For instance, NaCl and FeCls speed up corrosion, but a small amount of K2Cr207will greatly overcome their effects. In another experiment, obtain a few steel ball bearings, new or used, and apply sufficient pressure to the balls to make flat areas of about half the ball's diameter. This requires considerable pressure such as is available only in machine tool shops. Now, drop sueh a hearing into dilute HCI and measure the time far evidence of abrupt change-an audible rupture of the bearing or its obvious cleavage into hemispheres. Repeat this test by adding to the HCI some inhibitor, such as K2CnOl, methylmarpholine, or any typical commercial radiator rust inhibitor. and note the time for the hearing to fall in two. This experiment should be performed with moderate safety protection because occasionally the bursting ball bearing breaks the beaker in which the test is carried out. The test with prestressed ball bearings, incidentally, is a rapid method of screening and evaluating the effectiveness of corrosion inhibitors. It is probably based on the rate of stress produced by H2 gas at cathodic areas which, added to the stress already present in the bearing, causes the bearing to rupture. This is similar to so-called hydrogen emhrittlement. A good reference to corrosion chemistry is H. H. Uhlig, "Corrosion and Corrosion Control," John Wiley and Sons, Inc., New York, 1963.
