1.3 Kinetic Principles Thermodynamic principles can help explain a corrosion situation in terms of the stability of chemical species and reactions associated with corrosion processes. However, thermodynamic calculations cannot be used to predict corrosion rates. When two metals are put in contact, they can produce a voltage, as in a battery or electrochemical cell (see Galvanic Corrosion in Sec. 5.2.1). The material lower in what has been called the “galvanic series” will tend to become the anode and corrode, while the material higher in the series will tend to support a cathodic reaction. Iron or aluminum, for example, will have a tendency to cor- rode when connected to graphite or platinum. What the series cannot predict is the rate at which these metals corrode. Electrode kinetic principles have to be used to estimate these rates. 1.3.1 Kinetics at equilibrium: the exchange current concept The exchange current I 0 is a fundamental characteristic of electrode behavior that can be defined as the rate of oxidation or reduction at an equilibrium electrode expressed in terms of current. The term exchange current, in fact, is a misnomer, since there is no net current flow. It is merely a convenient way of representing the rates of oxida- tion and reduction of a given single electrode at equilibrium, when no loss or gain is experienced by the electrode material. For the corrosion of iron, Eq. (1.1), for example, this would imply that the exchange cur- 32 Chapter One Stresses due to corrosion product buildup Voluminous corrosion products Cracking and spalling of the concrete cover Reinforcing steel Reduced pH levels due to carbonation Figure 1.15 Graphical representation of the corrosion of reinforcing steel in concrete leading to cracking and spalling. 0765162_Ch01_Roberge 9/1/99 2:46 Page 32 rent is related to the current in each direction of a reversible reaction, i.e., an anodic current I a representing Eq. (1.7) and a cathodic current I c representing Eq. (1.8). Fe → Fe 2ϩ ϩ 2e Ϫ (1.7) Fe ← Fe 2ϩ ϩ 2e Ϫ (1.8) Since the net current is zero at equilibrium, this implies that the sum of these two currents is zero, as in Eq. (1.9). Since I a is, by con- vention, always positive, it follows that, when no external voltage or current is applied to the system, the exchange current is as given by Eq. (1.10). I a ϩ I c ϭ 0 (1.9) I a ϭϪI c ϭ I 0 (1.10) There is no theoretical way of accurately determining the exchange current for any given system. This must be determined experimental- ly. For the characterization of electrochemical processes, it is always preferable to normalize the value of the current by the surface area of the electrode and use the current density, often expressed as a small i, i.e., i ϭ I/surface area. The magnitude of exchange current density is a function of the following main variables: 1. Electrode composition. Exchange current density depends upon the composition of the electrode and the solution (Table 1.1). For redox reactions, the exchange current density would depend on the composi- tion of the electrode supporting an equilibrium reaction (Table 1.2). Aqueous Corrosion 33 TABLE 1.1 Exchange Current Density (i 0 ) for M z+ /M Equilibrium in Different Acidified Solutions (1M) Electrode Solution log 10 i 0 , A/cm 2 Antimony Chloride Ϫ4.7 Bismuth Chloride Ϫ1.7 Copper Sulfate Ϫ4.4; Ϫ1.7 Iron Sulfate Ϫ8.0; Ϫ8.5 Lead Perchlorate Ϫ3.1 Nickel Sulfate Ϫ8.7; Ϫ6.0 Silver Perchlorate 0.0 Tin Chloride Ϫ2.7 Titanium Perchlorate Ϫ3.0 Titanium Sulfate Ϫ8.7 Zinc Chloride Ϫ3.5; Ϫ0.16 Zinc Perchlorate Ϫ7.5 Zinc Sulfate Ϫ4.5 0765162_Ch01_Roberge 9/1/99 2:46 Page 33 Table 1.3 contains the approximate exchange current density for the reduction of hydrogen ions on a range of materials. Note that the val- ue for the exchange current density of hydrogen evolution on platinum is approximately 10 Ϫ2 A/cm 2 , whereas that on mercury is 10 Ϫ13 A/cm 2 . 2. Surface roughness. Exchange current density is usually expressed in terms of projected or geometric surface area and depends upon the surface roughness. The higher exchange current density for the H ϩ /H 2 system equilibrium on platinized platinum (10 Ϫ2 A/cm 2 ) compared to that on bright platinum (10 Ϫ3 A/cm 2 ) is a result of the larg- er specific surface area of the former. 3. Soluble species concentration. The exchange current is also a complex function of the concentration of both the reactants and the products involved in the specific reaction described by the exchange current. This function is particularly dependent on the shape of the charge transfer barrier ␤ across the electrochemical interface. 34 Chapter One TABLE 1.2 Exchange Current Density (i 0 ) at 25°C for Some Redox Reactions System Electrode Material Solution log 10 i 0 , A/cm 2 Cr 3ϩ /Cr 2ϩ Mercury KCl Ϫ6.0 Ce 4ϩ /Ce 3ϩ Platinum H 2 SO 4 Ϫ4.4 Fe 3ϩ /Fe 2ϩ Platinum H 2 SO 4 Ϫ2.6 Rhodium H 2 SO 4 Ϫ7.8 Iridium H 2 SO 4 Ϫ2.8 Palladium H 2 SO 4 Ϫ2.2 H ϩ /H 2 Gold H 2 SO 4 Ϫ3.6 Lead H 2 SO 4 Ϫ11.3 Mercury H 2 SO 4 Ϫ12.1 Nickel H 2 SO 4 Ϫ5.2 Tungsten H 2 SO 4 Ϫ5.9 O 2 reduction Platinum Perchloric acid Ϫ9.0 Platinum 10%–Rhodium Perchloric acid Ϫ9.0 Rhodium Perchloric acid Ϫ8.2 Iridium Perchloric acid Ϫ10.2 TABLE 1.3 Approximate Exchange Current Density (i 0 ) for the Hydrogen Oxidation Reaction on Different Metals at 25°C Metal log 10 i 0 , A/cm 2 Pb, Hg Ϫ13 Zn Ϫ11 Sn, Al, Be Ϫ10 Ni, Ag, Cu, Cd Ϫ7 Fe, Au, Mo Ϫ6 W, Co, Ta Ϫ5 Pd, Rh Ϫ4 Pt Ϫ2 0765162_Ch01_Roberge 9/1/99 2:46 Page 34 4. Surface impurities. Impurities adsorbed on the electrode sur- face usually affect its exchange current density. Exchange current den- sity for the H ϩ /H 2 system is markedly reduced by the presence of trace impurities like arsenic, sulfur, and antimony. 1.3.2 Kinetics under polarization When two complementary processes such as those illustrated in Fig. 1.1 occur over a single metallic surface, the potential of the material will no longer be at an equilibrium value. This deviation from equilib- rium potential is called polarization. Electrodes can also be polarized by the application of an external voltage or by the spontaneous pro- duction of a voltage away from equilibrium. The magnitude of polar- ization is usually measured in terms of overvoltage ␩, which is a measure of polarization with respect to the equilibrium potential E eq of an electrode. This polarization is said to be either anodic, when the anodic processes on the electrode are accelerated by changing the spec- imen potential in the positive (noble) direction, or cathodic, when the cathodic processes are accelerated by moving the potential in the neg- ative (active) direction. There are three distinct types of polarization in any electrochemical cell, the total polarization across an electro- chemical cell being the summation of the individual elements as expressed in Eq. (1.11): ␩ total ϭ␩ act ϩ␩ conc ϩ iR (1.11) where ␩ act ϭ activation overpotential, a complex function describing the charge transfer kinetics of the electrochemical processes. ␩ act is predominant at small polarization cur- rents or voltages. ␩ conc ϭ concentration overpotential, a function describing the mass transport limitations associated with electrochemi- cal processes. ␩ conc is predominant at large polarization currents or voltages. iR ϭ ohmic drop. iR follows Ohm’s law and describes the polar- ization that occurs when a current passes through an electrolyte or through any other interface, such as surface film, connectors, etc. Activation polarization. When some steps in a corrosion reaction con- trol the rate of charge or electron flow, the reaction is said to be under activation or charge-transfer control. The kinetics associated with apparently simple processes rarely occur in a single step. The overall anodic reaction expressed in Eq. (1.1) would indicate that metal atoms Aqueous Corrosion 35 0765162_Ch01_Roberge 9/1/99 2:46 Page 35 in the metal lattice are in equilibrium with an aqueous solution contain- ing Fe 2ϩ cations. The reality is much more complex, and one would need to use at least two intermediate species to describe this process, i.e., Fe lattice → Fe ϩ surface Fe ϩ surface → Fe 2ϩ surface Fe 2ϩ surface → Fe 2ϩ solution In addition, one would have to consider other parallel processes, such as the hydrolysis of the Fe 2ϩ cations to produce a precipitate or some other complex form of iron cations. Similarly, the equilibrium between protons and hydrogen gas [Eq. (1.2)] can be explained only by invoking at least three steps, i.e., H ϩ → H ads H ads ϩ H ads → H 2 (molecule) H 2 (molecule) → H 2 (gas) The anodic and cathodic sides of a reaction can be studied individual- ly by using some well-established electrochemical methods in which the response of a system to an applied polarization, current or voltage, is studied. A general representation of the polarization of an electrode sup- porting one redox system is given in the Butler-Volmer equation (1.12): i reaction ϭ i 0 Ά exp ΂ ␤ reaction ␩ reaction ΃ Ϫ exp ΄ Ϫ (1 Ϫ␤ reaction ) ␩ reaction ΅· (1.12) where i reaction ϭ anodic or cathodic current ␤ reaction ϭ charge transfer barrier or symmetry coefficient for the anodic or cathodic reaction, close to 0.5 ␩ reaction ϭ E applied Ϫ E eq , i.e., positive for anodic polarization and negative for cathodic polarization n ϭ number of participating electrons R ϭ gas constant T ϭ absolute temperature F ϭ Faraday nF ᎏ RT nF ᎏ RT 36 Chapter One 0765162_Ch01_Roberge 9/1/99 2:46 Page 36 When ␩ reaction is anodic (i.e., positive), the second term in the Butler- Volmer equation becomes negligible and i a can be more simply expressed by Eq. (1.13) and its logarithm, Eq. (1.14): i a ϭ i 0 ΄ exp ΂ ␤ a ␩ a ΃΅ (1.13) ␩ a ϭ b a log 10 ΂΃ (1.14) where b a is the Tafel coefficient that can be obtained from the slope of a plot of ␩ against log i, with the intercept yielding a value for i 0 . b a ϭ 2.303 (1.15) Similarly, when ␩ reaction is cathodic (i.e., negative), the first term in the Butler-Volmer equation becomes negligible and i c can be more sim- ply expressed by Eq. (1.16) and its logarithm, Eq. (1.17), with b c obtained by plotting ␩ versus log i [Eq. (1.18)]: i c ϭ i 0 Ά Ϫ exp ΄ Ϫ(1 Ϫ␤ c ) ␩ c ΅· (1.16) ␩ c ϭ b c log 10 ΂΃ (1.17) b c ϭϪ2.303 (1.18) Concentration polarization. When the cathodic reagent at the corroding surface is in short supply, the mass transport of this reagent could become rate controlling. A frequent case of this type of control occurs when the cathodic processes depend on the reduction of dissolved oxy- gen. Table 1.4 contains some data related to the solubility of oxygen in air-saturated water at different temperatures, and Table 1.5 contains some data on the solubility of oxygen in seawater of different salinity and chlorinity. 10 Because the rate of the cathodic reaction is proportional to the sur- face concentration of the reagent, the reaction rate will be limited by a drop in the surface concentration. For a sufficiently fast charge trans- fer, the surface concentration will fall to zero, and the corrosion process will be totally controlled by mass transport. As indicated in Fig. 1.16, mass transport to a surface is governed by three forces: dif- RT ᎏ ␤nF i c ᎏ i 0 nF ᎏ RT RT ᎏ ␤nF i a ᎏ i 0 nF ᎏ RT Aqueous Corrosion 37 0765162_Ch01_Roberge 9/1/99 2:46 Page 37 fusion, migration, and convection. In the absence of an electric field, the migration term is negligible, and the convection force disappears in stagnant conditions. For purely diffusion-controlled mass transport, the flux of a species O to a surface from the bulk is described with Fick’s first law (1.19), J O ϭϪD O ΂΃ (1.19) where J O ϭ flux of species O, mol и s Ϫ1 и cm Ϫ2 D O ϭ diffusion coefficient of species O, cm 2 и s Ϫ1 ϭ concentration gradient of species O across the interface, mol и cm Ϫ4 The diffusion coefficient of an ionic species at infinite dilution can be estimated with the help of the Nernst-Einstein equation (1.20), which relates D O to the conductivity of the species (␭ O ): ␦C O ᎏ ␦x ␦C O ᎏ ␦x 38 Chapter One TABLE 1.4 Solubility of Oxygen in Air-Saturated Water Temperature, °C Volume, cm 3 * Concentration, ppm Concentration (M), ␮mol/L 0 10.2 14.58 455.5 5 8.9 12.72 397.4 10 7.9 11.29 352.8 15 7.0 10.00 312.6 20 6.4 9.15 285.8 25 5.8 8.29 259.0 30 5.3 7.57 236.7 *cm 3 per kg of water at 0°C. TABLE 1.5 Oxygen Dissolved in Seawater in Equilibrium with a Normal Atmosphere Chlorinity,* % 0 5 10 15 20 Salinity,† % 0 9.06 18.08 27.11 36.11 Temperature, °C ppm 0 14.58 13.70 12.78 11.89 11.00 5 12.79 12.02 11.24 10.49 9.74 10 11.32 10.66 10.01 9.37 8.72 15 10.16 9.67 9.02 8.46 7.92 20 9.19 8.70 8.21 7.77 7.23 25 8.39 7.93 7.48 7.04 6.57 30 7.67 7.25 6.80 6.41 5.37 *Chlorinity refers to the total halogen ion content as titrated by the addition of silver nitrate, expressed in parts per thousand (%). †Salinity refers to the total proportion of salts in seawater, often estimated empirically as chlorinity ϫ 1.80655, also expressed in parts per thousand (%). 0765162_Ch01_Roberge 9/1/99 2:46 Page 38 D O ϭ (1.20) where z O ϭ the valency of species O R ϭ gas constant, i.e., 8.314 J и mol Ϫ1 и K Ϫ1 T ϭ absolute temperature, K F ϭ Faraday’s constant, i.e., 96,487 C и mol Ϫ1 Table 1.6 contains values for D O and ␭ O of some common ions. For more practical situations, the diffusion coefficient can be approximat- ed with the help of Eq. (1.21), which relates D O to the viscosity of the solution ␮ and absolute temperature: D O ϭ (1.21) where A is a constant for the system. TA ᎏ ␮ RT␭ O ᎏ |z O | 2 F 2 Aqueous Corrosion 39 H + e - H + H + 2e - e - H + Fe 2+ Fe 2+ Charge transfer Mass transport activation barrier ( )␣ exchange current density (i ) 0 Tafel slope (b) convection diffusion migration Figure 1.16 Graphical representation of the processes occurring at an electrochemical interface. 0765162_Ch01_Roberge 9/1/99 2:46 Page 39 TABLE 1.6 Conductivity and Diffusion Coefficients of Selected Ions at Infinite Dilution in Water at 25°C Cation |z| ␭, S и cm 2 и mol Ϫ1 D ϫ 10 5 , cm 2 и s Ϫ1 Anion |z| ␭, S и cm 2 и mol Ϫ1 D ϫ 10 5 , cm 2 и s Ϫ1 H ϩ 1 349.8 9.30 OH Ϫ 1 197.6 5.25 Li ϩ 1 38.7 1.03 F Ϫ 1 55.4 1.47 Na ϩ 1 50.1 1.33 Cl Ϫ 1 76.3 2.03 K ϩ 1 73.5 1.95 NO 3 Ϫ 1 71.4 1.90 Ca 2ϩ 2 119.0 0.79 ClO 4 Ϫ 1 67.3 1.79 Cu 2ϩ 2 107.2 0.71 SO 4 2Ϫ 2 160.0 1.06 Zn 2ϩ 2 105.6 0.70 CO 3 2Ϫ 2 138.6 0.92 O 2 —— 2.26 HSO 4 Ϫ 1 50.0 1.33 H 2 O —— 2.44 HCO 3 Ϫ1 1 41.5 1.11 40 0765162_Ch01_Roberge 9/1/99 2:46 Page 40 The region near the metallic surface where the concentration gra- dient occurs is also called the diffusion layer ␦. Since the concentra- tion gradient ␦C O /␦x is greatest when the surface concentration of species O is completely depleted at the surface (i.e., C O ϭ 0), it follows that the cathodic current is limited in that condition, as expressed by Eq. (1.22): i c ϭ i L ϭϪnFD O (1.22) For intermediate cases, ␩ conc can be evaluated using an expression [Eq. (1.23)] derived from the Nernst equation: ␩ conc ϭ log 10 ΂ 1 Ϫ ΃ (1.23) where 2.303RT/F ϭ 0.059 V when T ϭ 298.16 K. Ohmic drop. The ohmic resistance of a cell can be measured with a milliohmmeter by using a high-frequency signal with a four-point technique. Table 1.7 lists some typical values of water conductivity. 10 While the ohmic drop is an important parameter to consider when designing cathodic and anodic protection systems, it can be mini- mized, when carrying out electrochemical tests, by bringing the refer- ence electrode into close proximity with the surface being monitored. For naturally occurring corrosion, the ohmic drop will limit the influ- ence of an anodic or a cathodic site on adjacent metal areas to a cer- tain distance depending on the conductivity of the environment. For naturally occurring corrosion, the anodic and cathodic sites often are adjacent grains or microconstituents and the distances involved are very small. i ᎏ i L 2.303RT ᎏᎏ nF C O,, bulk ᎏ ␦ Aqueous Corrosion 41 TABLE 1.7 Resistivity of Waters Water ␳, ⍀иcm Pure water 20,000,000 Distilled water 500,000 Rainwater 20,000 Tap water 1000–5000 River water (brackish) 200 Seawater (coastal) 30 Seawater (open sea) 20–25 0765162_Ch01_Roberge 9/1/99 2:46 Page 41 [...]... Butterworth Heinemann, 19 94 076 516 2_ Ch 01_ Roberge 9 /1/ 99 2: 46 Page 13 Chapter 1 Aqueous Corrosion 1. 1 Introduction 13 1. 2 Applications of Potential-pH Diagrams 16 1. 2 .1 Corrosion of steel in water at elevated temperatures 17 1. 2. 2 Filiform corrosion 26 1. 2. 3 Corrosion of reinforcing steel in concrete 1. 3 Kinetic Principles 29 32 1. 3 .1 Kinetics at equilibrium: the exchange current concept 32 1. 3 .2 Kinetics under... i0 ϭ 10 Ϫ6 A и cm 2 I0 ϭ 1 ϫ 10 Ϫ6 A ba ϭ 0 . 12 0 V/decade Cathodic reaction Surface area ϭ 1 cm2 2Hϩ ϩ 2eϪ → H2 Eeq ϭ E0 ϩ 0.059 log10aHϩ ϭ 0.0 Ϫ 0.059 ϫ (Ϫ5) ϭ Ϫ0 .29 5 V versus SHE i0 ϭ 10 Ϫ6 A и cm 2 I0 ϭ 1 10 Ϫ6 A bc ϭ Ϫ0 . 12 0 V/decade 076 516 2_ Ch 01_ Roberge 9 /1/ 99 2: 46 Page 48 -0 .2 -0.3 Zn → Zn2+ + 2e- Potential (V vs SHE) -0.4 -0.5 -0.6 -0.7 -0.8 Ecorr & Icorr 2H++ 2e- → H2 -0.9 -1 -7 -6 -5 -4 -3 -2 -1 Log... I1 ϭ Surface area ϭ 1 cm2 076 516 2_ Ch 01_ Roberge 9 /1/ 99 2: 46 1 Page 51 O2 + 4H++ 4e- → 2H 2O 0.8 Potential (V vs SHE) 0.6 0.4 0 .2 Ecorr & Icorr 2H++ 2e- → H2 0 -0 .2 -0.4 Fe → Fe2+ + 2e- -0.6 -0.8 -8 -7 -6 -5 -4 -3 -2 Log (I(A)) Figure 1. 23 The iron mixed-potential diagram at 25 °C and pH 5 in an aerated solution with a limiting current of 10 Ϫ4 A for the reduction of oxygen 0 O2 + 4H++ 4e- → 2H2O -0 .1. .. solution at 25 °C, pH ϭ 5, Anodic reaction Surface area ϭ 1 cm2 Fe → Fe2ϩ ϩ 2eϪ For a corroding metal, one can assume that Eeq ϭ E0 i0 ϭ 10 Ϫ6 A и cm 2 I0 ϭ 1 ϫ 10 Ϫ6 A ba ϭ 0 . 12 0 V/decade Cathodic reactions Surface area ϭ 1 cm2 2Hϩ ϩ 2eϪ → H2 Eeq ϭ E0 ϩ 0.059 log10aHϩ ϭ 0.0 ϩ 0.059 ϫ (Ϫ5) ϭ Ϫ0 .29 5 V versus SHE i0 ϭ 10 Ϫ6 A и cm 2 I0 ϭ 1 ϫ 10 Ϫ6 A bc ϭ Ϫ0 . 12 0 V/decade O2 ϩ 4Hϩ ϩ 4eϪ → 2H2O E0 ϭ 1. 22 9 V versus... Figure 1. 26 The polarization curve corresponding to iron in a pH 2 solution at 25 °C in an aerated solution with a limiting current of 10 Ϫ4.5 A for the reduction of oxygen (Fig 1. 25 ) 52 076 516 2_ Ch 01_ Roberge 9 /1/ 99 2: 46 1 Page 53 O2 + 4H++ 4e- → 2H2O 0.8 Potential (V vs SHE) 0.6 0.4 0 .2 Ecorr & Icorr 0 2H++ 2e- → H2 -0 .2 Fe → Fe2+ + 2e- -0.4 -0.6 -0.8 -8 -7 -6 -5 -4 -3 -2 Log (I(A)) Figure 1. 27 The iron... Fe(OH) ;;;;;; HFeO ;;;;;;;;;;;;;;; ;;;;;; ;;;;;; ;;;;;; 2 6 4 8 10 12 pH Potential (V vs SHE) 1. 6 0.8 Recommended pH operating range to minimize corrosion damage 3 0 2+ 2 2 -0.8 Fe -1. 6 0 2 4 6 8 10 12 14 pH Figure 1. 5 E-pH diagram of iron in water at 25 °C and its observed corrosion behavior 076 516 2_ Ch 01_ Roberge 9 /1/ 99 2: 46 Page 21 Aqueous Corrosion 21 Potential (V vs SHE) example, no fewer than three... log10aHϩ ϩ (0.059/4) log10(pO2) Supposing pO 2 ϭ 0 .2, Eeq ϭ 1. 22 9 Ϫ 0.059 ϫ (Ϫ5) ϩ 0. 014 8 ϫ (Ϫ0.699) ϭ 0. 923 7 V versus SHE i0 ϭ 10 Ϫ7 A и cm 2 I0 ϭ 1 ϫ 10 Ϫ7 A bc ϭ Ϫ0 . 12 0 V/decade i1 ϭ I1 ϭ 10 Ϫ4 A The mixed-potential diagram of this system is shown in Fig 1. 23 , and the resultant polarization plot of the system is shown in Fig 1. 24 Fifth case: 10 Ϫ4.5 A iron in an aerated neutral solution at 25 °C, pH ϭ 2, ... a deaerated neutral solution at 25 °C, pH ϭ 5 076 516 2_ Ch 01_ Roberge 9 /1/ 99 2: 46 Page 47 Aqueous Corrosion 47 0 .2 + Potential (V vs SHE) - 2H + 2e → H 2 0 -0 .2 -0.4 E corr & I corr -0.6 Zn → Zn -0.8 -1 - 12 -11 -10 -9 2+ -8 + 2e - -7 -6 -5 -4 -3 -2 Log (I(A)) Figure 1. 19 The zinc mixed-potential diagram at 25 °C and pH 0 Anodic reaction Surface area ϭ 1 cm2 Fe → Fe2ϩ ϩ 2eϪ E0 ϭ Ϫ0.44 V versus SHE For a... → Fe2+ + 2e-0 .2 -0.3 Ecorr & Icorr -0.4 -0.5 -0.6 2H++ 2e- → H2 -0.7 -6 -5.5 -5 -4.5 -4 -3.5 -3 -2. 5 -2 Log (I(A)) Figure 1. 24 The polarization curve corresponding to iron in a pH 5 solution at 25 °C in an aerated solution with a limiting current of 10 Ϫ4 A for the reduction of oxygen (Fig 1. 23 ) 51 076 516 2_ Ch 01_ Roberge 9 /1/ 99 2: 46 1 Page 52 O2 + 4H++ 4e- → 2H2O 0.8 Potential (V vs SHE) 0.6 0.4 0 .2 Ecorr... Log (I(A)) Figure 1. 20 The polarization curve corresponding to zinc in a pH 0 solution at 25 °C (Fig 1. 19) 0 Potential (V vs SHE) -0 .1 -0 .2 2H++ 2e- → H2 -0.3 Ecorr & Icorr -0.4 Fe → Fe2+ + 2e-0.5 -0.6 -8 -7 -6 -5 -4 Log (I(A)) Figure 1. 21 The iron mixed-potential diagram at 25 °C and pH 5 48 -3 -2 076 516 2_ Ch 01_ Roberge 9 /1/ 99 2: 46 Page 49 Aqueous Corrosion 49 -0 .2 Fe → Fe2+ + 2e-0 .25 Potential (V vs . 10 15 20 Salinity,† % 0 9.06 18 .08 27 .11 36 .11 Temperature, °C ppm 0 14 .58 13 .70 12 .78 11 .89 11 .00 5 12 .79 12 . 02 11 .24 10 .49 9.74 10 11 . 32 10 .66 10 . 01 9.37 8. 72 15 10 .16 9.67 9. 02 8.46 7. 92 20. 50 .1 1.33 Cl Ϫ 1 76.3 2. 03 K ϩ 1 73.5 1. 95 NO 3 Ϫ 1 71. 4 1. 90 Ca 2 2 11 9.0 0.79 ClO 4 Ϫ 1 67.3 1. 79 Cu 2 2 10 7 .2 0. 71 SO 4 2 2 16 0.0 1. 06 Zn 2 2 10 5.6 0.70 CO 3 2 2 13 8.6 0. 92 O 2 —— 2. 26. (M), ␮mol/L 0 10 .2 14 .58 455.5 5 8.9 12 . 72 397.4 10 7.9 11 .29 3 52. 8 15 7.0 10 .00 3 12 .6 20 6.4 9 .15 28 5.8 25 5.8 8 .29 25 9.0 30 5.3 7.57 23 6.7 *cm 3 per kg of water at 0°C. TABLE 1. 5 Oxygen Dissolved