Handbook of Corrosion Engineering Episode 1 Part 3 pot

alloying additions or protective coatings for corrosion resistance are associated with this steel. In simplistic terms, concrete is produced by mixing cement clinker, water, fine aggregate (sand), coarse aggregate (stone), and other chem- ical additives. When mixed with water, the anhydrous cement clinker compounds hydrate to form cement paste. It is the cement paste that forms the matrix of the composite concrete material and gives it its strength and rigidity, by means of an interconnected network in which the aggregate particles are embedded. The cement paste is porous in nature. An important feature of concrete is that the pores are filled with a highly alkaline solution, with a pH between 12.6 and 13.8 at normal humidity levels. This highly alkaline pore solution arises from by-products of the cement clinker hydration reactions such as NaOH, KOH, and Ca(OH) 2 . The maintenance of a high pH in the concrete pore solution is a fundamental feature of the corrosion resistance of carbon steel reinforcing bars. At the high pH levels of the concrete pore solution, without the ingress of corrosive species, reinforcing steel embedded in concrete tends to display completely passive behavior as a result of the forma- tion of a thin protective passive film. The corrosion potential of passive reinforcing steel tends to be more positive than about Ϫ0.52 V (SHE) according to ASTM guidelines. 9 The E-pH diagram in Fig. 1.14 con- firms the passive nature of steel under these conditions. It also indi- cates that the oxygen reduction reaction is the cathodic half-cell reaction applicable under these highly alkaline conditions. One mechanism responsible for severe corrosion damage to reinforc- ing steel is known as carbonation. In this process, carbon dioxide from the atmosphere reacts with calcium hydroxide (and other hydroxides) in the cement paste following reaction (1.6). Ca(OH) 2 ϩ CO 2 → CaCO 3 ϩ H 2 O (1.6) The pore solution is effectively neutralized by this reaction. Carbonation damage usually appears as a well-defined “front” parallel to the outside surface. Behind the front, where all the calcium hydrox- ide has reacted, the pH is reduced to around 8, whereas ahead of the front, the pH remains above 12.6. When the carbonation front reaches the reinforcement, the passive film is no longer stable, and active cor- rosion is initiated. Figure 1.14 shows that active corrosion is possible at the reduced pH level. Damage to the concrete from carbonation- induced corrosion is manifested in the form of surface spalling, result- ing from the buildup of voluminous corrosion products at the concrete-rebar interface (Fig. 1.15). A methodology known as re-alkalization has been proposed as a remedial measure for carbonation-induced reinforcing steel corro- 30 Chapter One 0765162_Ch01_Roberge 9/1/99 2:46 Page 30 sion. The aim of this treatment is to restore alkalinity around the reinforcing bars of previously carbonated concrete. A direct current is applied between the reinforcing steel cathode and external anodes positioned against the external concrete surface and surrounded by electrolyte. Sodium carbonate has been used as the electrolyte in this process, which typically requires several days for effectiveness. Potential disadvantages of the treatment include reduced bond strength, increased risk of alkali-aggregate reaction, microstructural changes in the concrete, and hydrogen embrittlement of the reinforc- ing steel. It is apparent from Fig. 1.14 that hydrogen reduction can occur on the reinforcing steel cathode if its potential drops to highly negative values. Aqueous Corrosion 31 pH Potential (V vs SHE) 1.6 0.8 0 -0.8 -1.6 0 2 46 810 12 14 A B Fe Fe 2+ Decreasing pH from carbonation makes shift to active field possible Potential range associated with passive reinforcing steel Re-alkalization attempts to re-establish passivity HFeO 2 - Fe O 34 Figure 1.14 E-pH diagram of the iron-water system with an emphasis on the microenviron- ments produced during corrosion of reinforcing steel in concrete. 0765162_Ch01_Roberge 9/1/99 2:46 Page 31 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 [...]...076 516 2_Ch 01_ Roberge 40 Conductivity and Diffusion Coefficients of Selected Ions at Infinite Dilution in Water at 25°C |z| ␭, S и cm2 и mol 1 Hϩ 1 34 9.8 Liϩ 1 38 .7 Naϩ Kϩ 1 1 50 .1 73. 5 D ϫ 10 5, cm2 и s 1 Anion |z| ␭, S и cm2 и mol 1 D ϫ 10 5, cm2 и s 1 9 .30 OHϪ 1 197.6 5.25 1. 03 FϪ 1 55.4 1. 47 1. 33 ClϪ 1 76 .3 2. 03 1. 95 NO3Ϫ 1 71. 4 1. 90 Ϫ Ca2ϩ 2 11 9.0 0.79 ClO4 1 67 .3 1. 79 Cu2ϩ 2 10 7.2 0. 71 SO42Ϫ 2 16 0.0... 12 3 12 4 12 6 12 9 12 9 Salinity 12 9 Other ions 13 1 Precipitation of inorganic compounds from seawater 13 1 Oxygen 13 3 Organic compounds 13 5 Polluted seawater 13 6 Brackish coastal water 13 7 2 .3. 2 Corrosion resistance of materials in seawater 13 8 Carbon steel 13 9 Stainless steels 14 0 Nickel-based alloys 14 0 Copper-based alloys 14 0 Effect of flow velocity 14 0 Effect of temperature 2.4 Corrosion in Soils 14 1... Soils 14 1 14 2 2.4 .1 Introduction 14 2 2.4.2 Soil classification systems 14 2 2.4 .3 Soil parameters affecting corrosivity 14 3 Water 14 3 Degree of aeration 14 3 pH 14 3 Soil resistivity 14 6 Redox potential 14 6 Chlorides 14 6 Sulfates 14 7 Microbiologically influenced corrosion 2.4.4 Soil corrosivity classifications 2.4.5 Corrosion characteristics of selected metals and alloys Ferrous alloys 14 7 14 8 15 1 15 1 Nonferrous... 1. 06 Zn2ϩ 2 10 5.6 0.70 CO32Ϫ 2 13 8.6 0.92 2.26 HSO4Ϫ 1 50.0 1. 33 2.44 HCO3 1 1 41. 5 1. 11 O2 H2O — — — — Page 40 Cation 9 /1/ 99 2:46 TABLE 1. 6 076 516 2_Ch 01_ Roberge 9 /1/ 99 2:46 Page 41 Aqueous Corrosion 41 The region near the metallic surface where the concentration gradient occurs is also called the diffusion layer ␦ Since the concentration gradient ␦CO/␦x is greatest when the surface concentration of. .. Eeq ϭ 1. 229 Ϫ 0.059 ϫ (Ϫ5) ϩ 0. 014 8 ϫ (Ϫ0.699) ϭ 0.9 237 V versus SHE i0 ϭ 10 Ϫ7 A и cmϪ2 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, I1 ϭ Surface area ϭ 1 cm2 076 516 2_Ch 01_ Roberge 9 /1/ 99... resins 10 1 Selectivity of resins 10 3 Kinetics 10 3 Types of ion-exchange resins 2.2 .3 Saturation and scaling indices 10 4 10 5 The Langelier saturation index 10 6 Ryznar stability index 10 8 Puckorius scaling index 10 8 Larson-Skold index 10 9 55 076 516 2_Ch02_Roberge 56 9 /1/ 99 4: 01 Page 56 Chapter Two Stiff-Davis index 11 0 Oddo-Tomson index 11 0 Momentary excess (precipitation to equilibrium) 11 0 Interpreting the... and alloys 15 1 Reinforced concrete 15 3 2.4.6 Summary 15 4 076 516 2_Ch02_Roberge 9 /1/ 99 4: 01 Page 57 Environments 2.5 Reinforced Concrete 57 15 4 2.5 .1 Introduction 15 4 2.5.2 Concrete as a structural material 15 5 2.5 .3 Corrosion damage in reinforced concrete 15 6 Mehta’s holistic model of concrete degradation 15 6 Corrosion mechanisms 15 9 Chloride-induced rebar corrosion 15 9 Carbonation-induced corrosion. .. 11 1 11 2 Optimizing storage conditions for low-level nuclear waste 11 4 Limiting halite deposition in a wet high-temperature gas well 11 5 Identifying acceptable operating range for ozonated cooling systems 11 7 Optimizing calcium phosphate scale inhibitor dosage in a high-TDS cooling system 12 2 2.2.5 Software systems Scaling of cooling water Scaling of deep well water 2 .3 Seawater 2 .3 .1 Introduction 12 3. .. measures 16 5 16 6 Alternative deicing methods 16 6 Cathodic protection 16 8 Electrochemical chloride extraction 17 0 Re-alkalization 17 1 Repair techniques 17 3 Epoxy-coated reinforcing steel 17 5 Stainless steel rebar 17 5 Galvanized rebars 17 7 Corrosion inhibitors 17 8 Concrete cover and mix design 17 8 2.5.5 Condition assessment of reinforced concrete structures Electrochemical corrosion measurements 18 0 18 2 Chloride... content 18 3 Petrographic examination 18 4 Permeability tests 18 4 2.5.6 Life prediction for corroding reinforced concrete structures 2.5.7 Other forms of concrete degradation Alkali-aggregate reaction 18 4 18 6 18 6 Freeze-thaw damage 2.6 18 7 Sulfate attack 18 7 Microbes and Biofouling 2.6 .1 Basics of microbiology and MIC Classification of microorganisms 18 7 18 7 19 0 Bacteria commonly associated with MIC 19 1 Effect . 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. mol 1 D ϫ 10 5 , cm 2 и s 1 Anion |z| ␭, S и cm 2 и mol 1 D ϫ 10 5 , cm 2 и s 1 H ϩ 1 34 9.8 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. 2. 03 K ϩ 1 73. 5 1. 95 NO 3 Ϫ 1 71. 4 1. 90 Ca 2ϩ 2 11 9.0 0.79 ClO 4 Ϫ 1 67 .3 1. 79 Cu 2ϩ 2 10 7.2 0. 71 SO 4 2Ϫ 2 16 0.0 1. 06 Zn 2ϩ 2 10 5.6 0.70 CO 3 2Ϫ 2 13 8.6 0.92 O 2 —— 2.26 HSO 4 Ϫ 1 50.0 1. 33 H 2 O ——