1 Introduction Corrosion is the unwanted reaction or destruction of a metal component by the environment. The annual cost of corrosion to the US economy has been estimated to be over $70 billion. Similar costs are associated with other industrialized countries. Many of the problems can be avoided if basic precautions and design processes are followed.
The mechanism of corrosion is electrochemical and can be induced by the flow of current or will cause a current to flow. When a corroding metal is oxidized, the reaction M ! Mþ þ ne n (10.1) must be accompanied by a reduction reaction which is usually the reduction of oxygen whether in the air or dissolved in water. Gesser, Applied Chemistry: A Textbook for Engineers and Technologists, 175 DOI 10.1007/978-1-4614-4262-2_10, # Springer Science+Business Media New York 2013 176 10 Corrosion Or O2 þ 2H2 O þ 4e ! 4OH (10.3) In some cases, the reduction of hydrogen occurs.4) The usual classification of corrosion is according to the environment to which the metal is exposed or the actual reactions which occur. We have seen that the concentration cell is a simple cell in which a metal can corrode as dissolution takes place.2 Factors Affecting the Rate of Corrosion It is convenient to classify the corrosion of metals in terms of (a) the metals and (b) the environment.
The reduction potential is the most important characteristic of a metal that determines its susceptibility to corrosion. This has been illustrated by Table 9. Thus, the noble metals, gold and platinum, are resistant to corrosion and will only dissolve in strong oxidizing solutions which also contain complexing halides or other ions, for example, (CN–). For metals in seawater, the relative order of the reduction potential of metals and alloys has been established.
This is illustrated in Table 10.1 where distinction is made between active and passive surfaces for some metals. Magne- sium is a most active metal, whereas platinum and graphite are the least active materials. The voltages are given with respect to the saturated calomel electrode (SCE).1 The oxidation reaction (10.1) represents corrosion which must be accompanied by a reduction reaction (10.4) as well as reactions such as þ Fe3þ þ e ! Fe2 (10.5) and 3Hþ þ NO3 þ 2e ! HNO2 þ H2 O (10.6) The reaction which occurs depends on the solution in which the metal corrodes, but in most cases the cathodic reaction involves O2. The corrosion rate will thus depend on the partial pressure of oxygen.
This is shown in Table 10. Hence, the removal of oxygen from water in steam boilers is one method of reducing corrosion. If hydrogen evolution is the cathodic reaction (10.4), then it can be reduced by increasing the overvoltage. The overvoltage of H2 on mercury is very high (see Table 9.4) can be inhibited if mercury is used to coat the metal surface and to form an amalgam (see the zinc–air cell, Sect.
The overvoltage is dependent on current density which is determined by the area of the metals. Hence, as the cathode area decreases, the polarization can be expected to increase resulting in a decrease in rate of corrosion. In the case of iron (anode) on a large copper sheet (cathode), the large cathode/anode ratio favors corrosion of the iron. This is shown in Fig.
1 The saturated calomel electrode is a convenient reference electrode often used instead of the standard hydrogen electrode: 12 Hg2 Cl2 þ e ! Hg þ Cl , ℰ ¼ 0.2 Factors Affecting the Rate of Corrosion 177 Table 10.1 Galvanic metal V (V) and alloy potential V (vs. SCE) in seawater Mg 1.05 Low alloy steel 0.03 Solder Pb–Sn 0.02 Ni/Al bronze 0.05 Ni/Co 600 alloy 0.06 Ag bronze alloys 0.05 Alloy 2C, stainless steel 0.06 Ni/Fe/Cr/Alloy 825 0.04 Ni/Cr/Mo/Cu/Si alloy 0.06 Ni/Cr/Mo alloy C 0.16 The type and amount of impurities in a metal will affect the rate of corrosion. For example, a zinc sample which is 99.99 % pure (referred to as 4n zinc) would corrode about 2,000 times faster than a 5n sample. Even improperly annealed metals will show excessive corrosion rates.
Another factor which controls the rate of corrosion is the relative volume of the corrosion product (oxide) to the metal as well as the porosity of the oxide layer. For example, the volume ratio of oxide/ metal for Al, Ni, Cr, and W is 1. The oxide layer on a metal can 178 10 Corrosion Table 10.2 Effect of O2 P(O2) (atm) Rate of corrosion (mm/year) pressure on corrosion of iron in seawater 0.1 The corrosion of an iron rivet in a copper plate. The large copper surface results in a low O2 overvoltage, allowing the corrosion to proceed at a rate controlled by O2 diffusion convert a metal from one that corrodes to one that is inert.
Aluminum can react with water to form hydrogen by the reaction 2A1 þ 6H2 O ! 2AlðOHÞ3 þ 3H2 (10.7) followed by 2Al ðOHÞ3 ! Al2 O3 þ 3H2 O (10.8) However, the oxide layer which forms prevents the water from contacting the aluminum surface. Only in acid or alkali is the Al2O3 solubilized, and the aluminum reacts to liberate hydrogen. An oxide layer is readily formed on many metals when they are made anodic in aqueous solutions. In the case of aluminum, this process is called anodization.
It is also referred to as a passive film which reduces the corrosion rate. Such passive films can be thin, from 0.01 mm, and fragile and easily broken. Thus, when steel is immersed in nitric acid or chromic acid and then washed, the steel does not immediately tarnish nor will it displace copper from aqueous CuSO4. The steel has become passive due to the formation of an adhering oxide film which can be readily destroyed by HCl which forms the strong acid H+ FeCl4.
The factors influencing the rusting of iron can be illustrated by the electrochemical treatment of the overall reaction. þ 2Fe þ O2 þ 4Hþ ! 2Fe2 þ 2H2 O þ Fe ! Fe2 þ 2e E ¼ 0:440 V O2 þ 4Hþ þ 4e ! 2H2 O E ¼ 1:67 V 10.3 Types of Corrosion 179 From the Nernst equation (9.12) 2 ½Fe2þ E cell ¼ E 0 cell ðRT=4FÞ ln (10.9) Po2 ½Hþ 4 The corrosion reaction (10.9) ceases when ℰcell 0 2 ½Fe2þ 0 ¼ 1:67 ð0:0591=4Þ ln (10.10) Po2 ½Hþ 4 Hence, ℰcell 0 when 2 ½Fe2þ log 113 (10.11) Po2 ½Hþ 4 Let us consider extreme conditions where Po2 ¼ 106 atm; Fe2þ ¼ 10 M; ½Hþ ¼ 1014 M log 102 =ð106 1056 Þ ¼ log 1064 ¼ 64 Since 64 < 113, corrosion will continue to occur. In strong NaOH solution, rusting is reduced because the Fe2O3 forms a protective layer over the metal.3 Types of Corrosion The various forms of corrosion can be classified by their various causes. These are uniform corrosion attack (UC), bimetallic corrosion (BC), crevice corrosion (CC), pitting corrosion (PC), grain bound- ary corrosion (GBC), layer corrosion (LC), stress corrosion cracking (SCC), cavitation corrosion (CC), and hydrogen embrittlement (HE).1 Uniform Corrosion Such corrosion is usually easy to detect and rectify.
The slow corrosion of a metal in aqueous acidic solution is an example of such corrosion. Impurities in a metal can result in local cells which, in the presence of electrolyte, will show corrosive action.2 Bimetallic Corrosion This type of corrosion, also called galvanic corrosion, is characterized by the rapid dissolution of a more reactive metal in contact with a less reactive more noble metal. For example, galvanized steel (Zn–Fe) in contact with copper (Cu) pipe is a common household error. A nonconducting plastic 180 10 Corrosion Fig.2 A gradient in O2 concentration in the water drop makes the center portion of the iron anodic where Fe Fe- + 2e~, while the edge is cathodic and oxygen is reduced spacer would reduce the corrosion rate in the pipe.
The rate of corrosion is partially determined by the difference in the standard cell potentials of the two metals in contact (see Table 9. The relative potential of metals in seawater is given in Table 10.1 and represents the driving force of the corrosion which includes the current, or more precisely, the current density, that is, A/cm2. An electrochemical cell is formed and the anodic metal dissolves. This can be corrected by applying a counter current or voltage or by introducing a more reactive, sacrificial anode, for example, adding magnesium alloy to the above Zn–Fe–Cu system, a procedure commonly used for hot-water pipes in renovated buildings.3 Crevice Corrosion A nonuniform environment or concentration gradient due to material structure or design leads to concentration cells and corrosion.
Differential aeration is, for example, the cause of corrosion at the waterline or at the edges of holes or flange joints. The size of a crevice can range from 25 to 100 mm in width—small enough to create an oxygen concentration cell between the crevice solution and that on the outer surface. Oxygen can form a thin oxide layer on metals which acts as a protective passive film.4 Pitting Corrosion Like CC, PC is due to differential aeration or film formation (due to dust particles). The breakdown of a protective oxide layer at a lattice defects is another common cause of pits.
The mechanism of pitting of iron under a water drop is shown in Fig.2, and as in a CC, a differential concentration of oxygen in the drop creates a concentration cell. Rust has the composition of Fe3O4 and FeO(OH). Fe3O4 is a mixed oxide of FeO · Fe2O3 where iron is in the +2 and +3 oxidation state. The PC of various iron alloys induced by Cl– in the presence of 0.5 M H2SO4 is given in Table 10.
High chromium alloys are effective in reducing PC, but a limit is reached at about 25 % Cr, whereas nickel seems to have little effect on corrosion resistance. Other salts in solution also can affect the pitting rate as well as the depth of the pits.to the bulk alloy, and severe intergranular corrosion and pitting results. The corrosion rate of stainless steel (18/8) in aqueous HCl solutions depends on the concentration of acid, 10.3 Types of Corrosion 181 Table 10.3 Minimum Alloy [Cl–] (M) concentration of chloride ion necessary for starting pitting Fe 0.3 A comparison of the corrosion rates of metallic glasses (x, ) and crystalline stainless steel (O, A) as a function of HCl concentration at 30 C. No weight changes of the metallic glasses of Fe-Cr10P13C7 were detected by a microbalance after immersion for 200 hr temperature, and the oxygen pressure.
In contrast, an equivalent metallic glass2 (Fe–Cr10Ni5P13C7) showed no detectable corrosion. This is illustrated in Fig.3 and clearly shows how important corrosion is along the grain boundaries in stainless steel. Similar results were obtained for immersion tests in 10 % wt. of FeCl3 · 6H2O at 60 C as an indication of PC.
Again, the stainless steel (304, 136, 316) all showed significant pitting, whereas metallic glasses showed no detectable weight loss after 200 h. Not all metallic glasses are resistant to corrosion, and much more work is needed to understand these differences. 2 Metallic glasses are amorphous noncrystalline solids which are usually prepared by rapidly cooling the molten metal. Such metals are devoid of grain boundaries.5 Grain Boundary Corrosion Coarse crystalline-rolled metals or alloys can corrode at the edge of the crystallites; thus, the iron impurity in aluminum is responsible for aluminum corrosion.
Similarly, stainless steel (18/8 Cr/Ni) when heated (during welding) results in the precipitation of chromium carbide at the grain boundaries. This forms an enriched nickel layer anodic to the bulk alloy, and severe intergranular corrosion and pitting results. The corrosion of stainless steel (18/8) in aqueous HCl solutions depends on the concentration of acid, temperature, and the oxygen pressure.