Fig. 1 Three-electrode device. V, voltmeter E. Protopopoff and P. Marcus, Potential Measurements with Reference Electrodes, Corrosion: Fundamentals, Testing, and Protection, Vol 13A, ASM Handbook, ASM International, 2003, p 13–-16 Potential Measurements with Reference Electrodes E. Protopopoff, Laboratoire de Physico-Chimie des Surfaces, CNRS, and P. Marcus, Ecole Nationale Supérieure de Chimie de Paris, Université Pierre et Marie Curie Electrode Selection Characteristics A good reference electrode must reach its potential quickly, be reproducible, and remain stable with time. It must have a practically nonpolarizable metal-solution interface; that is, its potential must not depart significantly from the equilibrium value on the passage of a small current across the interface. The potential of the junction between the electrolytes of the reference and test electrodes must be minimized. These criteria are detailed subsequently. Stable and Reproducible Potential. Electrodes used as references should rapidly achieve a stable and reproducible potential that is free of significant fluctuations. To obtain these characteristics, it is advantageous, whenever possible, to use reversible electrodes, which can easily be made. The reference electrode arbitrarily chosen to establish a universal potential scale is the standard hydrogen electrode (SHE). It consists of a platinized or black platinum wire or sheet immersed in an aqueous solution of unit activity of protons saturated with hydrogen gas at a fugacity of one bar (14.5 psia). The half-cell reaction is H + (aq) + e - ↔ H 2 (g). Any non-standard reversible hydrogen electrode with well-controlled H + activity and H 2 fugacity can also be used as a reference. The equilibrium potential (E eq ) of a nonstandard reversible hydrogen electrode versus the SHE is, from the Nernst equation (Eq 25 of the article “Electrode Potentials” in this Section of the volume): (Eq 1) where R is the gas constant, T is the absolute temperature, F is the Faraday constant, is the proton activity in solution, and is the H 2 fugacity near the electrode; the SHE potential, , is, by convention, equal to zero. Platinization of a smooth platinum electrode is achieved by electrodeposition from a solution of H 2 PtCl 6 of a black platinum layer having a very rough surface and hence a very high specific area. The electrolyte is made up of a mineral acid (HCl or H 2 SO 4 ) with a well-defined activity of H + . Hydrogen gas is bubbled on the black platinum. While the hydrogen electrode is the fundamental reference, it has certain disadvantages in real conditions: the electrolyte must be prepared with an accurately known proton activity; the hydrogen gas must be purified, particularly from oxygen; furthermore, the platinum electrode must be frequently replatinized, because it easily gets “poisoned” by the adsorption of impurities present in the solution, which prevents the establishment of the equilibrium potential. For these reasons, practical corrosion measurements are usually not performed with the SHE but with secondary reference electrodes that are easier to construct and handle, less sensitive to impurities, and whose potential is very stable and well- known with respect to the SHE (Ref 1). The simpler reference electrodes are metal-ion (M z+ /M) electrodes, also called metallic electrodes of the first kind. The copper-copper sulfate (CuSO 4 /Cu) electrode is an excellent example of a good reversible electrode and it is widely used as a reference electrode in the corrosion field. It can easily be made by immersing a copper wire in a glass tube filled with a CuSO 4 aqueous solution and terminated by a porous plug (to allow ionic conduction with the cell electrolyte), as shown in Fig. 2. Fig. 2 Schematic of a copper/copper sulfate reference electrode This electrode is reversible, because a small cathodic current produces the reduction reaction (Cu 2+ + 2e - → Cu), while an anodic current brings about the oxidation reaction (Cu → Cu 2+ + 2e - ). Copper is a semi-noble metal and does not dissolve anodically in a solution of protons. In the case of the CuSO 4 /Cu electrode, the rest potential is equal to the equilibrium potential that can be computed from the Nernst equation: (Eq 2) where the standard potential = +0.337 V versus SHE (see Table 1 in the article “Electrode Potentials” in this Volume), and is the activity of Cu 2+ in the aqueous solution. If a copper solution of concentration 1.00 mol/L is used (where < 1), the equilibrium potential of the CuSO 4 /Cu electrode takes the value +0.310 V versus SHE at 25 °C (77 °F). This well-defined reversible electrode is reliable and easy to build. Another common metal-ion electrode is the silver electrode (Ag + /Ag). The Nernst equation applied to the half- cell reaction Ag + + e - ↔ Ag gives: (Eq 3) where the standard potential = 0.800 V versus SHE. Metallic electrodes of the second kind (also called secondary reference electrodes) are similar electrodes where the potential-determining activity of the metallic ion, , in solution is controlled by putting the metal M in contact with a sparingly soluble M compound (salt, oxide, hydroxide), itself in contact with the solution. Then, is determined by the solubility product of the salt and the activity of the anion. For example, if silver chloride (AgCl), only slightly soluble in water, is present on the silver surface, then the following equilibrium holds: AgCl ↔ Ag + + Cl - (Eq 4) The equilibrium mass law gives the solubility product as = K s(AgCl) . Expressing from this relation and placing it in Eq 3 leads to the following expression of the equilibrium potential of the Ag + , Cl - /AgCl/Ag half-cell: (Eq 5) corresponding to the global electrode equilibrium: AgCl + e - ↔ Ag + Cl - (Eq 6) The standard potential of this reversible silver-silver chloride (Ag/Cl) electrode is = + ln K s(AgCl) . At 25 °C, (77 °F), because = 0.800 V/SHE and K s(AgCl) = 1.78 · 10 -10 , = +0.222 V/SHE (Ref 2). From Eq 5, the potential of this type of reference electrode depends on the activity of the anion in solution (Cl - here), and this activity is controlled by the addition of a soluble salt of this anion (KCl). If a KCl solution of concentration 1.00 mol/L is used (where < 1), the equilibrium potential of the AgCl/Ag electrode takes the value +0.237 V/ SHE at 25 °C (77 °F). In a saturated KCl solution (where > 1), it is +0.198 V/SHE (Ref 2). Figure 3 shows a typical Ag-AgCl electrode. The AgCl coating is made by anodization of the silver wire in a chloride-containing solution. This electrode is easily assembled and can be placed directly into the electrochemical cell. Quite similar are the Pb/PbSO 4 and the Ag-Ag 2 O (Ref 1, 2). Fig. 3 Schematic of silver-silver chloride and calomel reference electrodes. Source: Ref 12 The most used secondary reference electrode is the calomel electrode (Hg-Hg 2 Cl 2 ). The sparingly soluble salt is, in this case, calomel (Hg 2 Cl 2 ) which floats as a paste on the top of a liquid mercury drop and which, in contact with a KCl solution, dissociates slightly into and Cl - ions. Using the same method as for the Ag- AgCl electrode, the following relation is obtained: (Eq 7) corresponding to the global electrode equilibrium: Hg 2 Cl 2 + 2e - ↔ 2Hg + 2Cl - (Eq 8) Figure 3 shows a typical calomel electrode. A platinum wire connects the electrode to the rest of the circuit. The most usual calomel electrodes are prepared with KCl solutions at a unit molarity of Cl - (normal calomel electrode or NCE) or saturated calomel electrode or (SCE). At 25 °C (77 °F), = 0.268 V/SHE, E NCE = 0.281 V, and E SCE = 0.242 V/SHE (Ref 2). The SCE has the advantage of being the easiest to prepare. However, due to the high temperature dependence of the KCl solubility, its potential varies markedly with temperature (~1 mV/°C as compared to 0.1 mV/°C for the NCE) (Ref 2), so its use makes mandatory the accurate monitoring of the temperature of experiment. Moreover, one must be careful when using a calomel electrode that only very low currents pass through the interface, because HgO may form, which irreversibly spoils the electrode. Similar reference electrodes for measurements in aqueous solutions are the mercurous sulfate electrode ( , ) in a solution of potassium sulfate, and the mercuric oxide electrode (Hg 2+ , OH - /HgO/Hg) in a solution of sodium hydroxide. Their potentials at 25 °C (77 °F) may be found in the literature (Ref 1, 2). All of these secondary reference electrodes are reversible. In some practical cases, nonreversible electrodes such as graphite are used. Although not as good, their potential stability in a particular environment is considered sufficient for certain applications. In the selection of reference electrodes, their durability and price must also be considered. Low Polarizability. A good reference electrode must have a practically nonpolarizable metal-solution interface; that is, its potential must not depart significantly from the equilibrium value on the passage across it of a small current (even minimized in the three-electrode device, Fig. 1), because the electrode polarization introduces an error in the potential measurement. The potential versus current density response of a good reference electrode, called a polarization curve, should be as flat as possible. Consider that the electrode equilibrium is bO + ne - = cR, where O is the oxidized species; R is the reduced species; and n is an integer. The net current density across the interface is proportional to the difference between the anodic and the cathodic rates. If the electrode reaction rate is limited by the electron transfer across the metal-solution interface (electron transfer is slow compared to mass transport of O and R between the bulk solution and the interface), the net current density is simply related to the overpotential η = E - E eq , where E is the potential, and E eq is the equilibrium potential, by (Ref 3): (Eq 9) where i 0 is the exchange current density of the electrode reaction; and α is the charge transfer coefficient for the anodic reaction (0 < α < 1), whose value is close to 0.5 for a single-step reaction (n = 1). (More detailed information on polarization curves can be found in the articles “Kinetics of Aqueous Corrosion” and “Electrochemical Methods of Corrosion Testing” in this Volume). For potentials close to the equilibrium potential, such as η < RT/F (26 mV at room temperature), the relation (Eq 9) can be approximated by a linear i versus η dependence: (Eq 10) or (Eq 11) The term RT/nFi 0 = (dη/di) 0 is called the polarization resistance, R p . A good reference electrode should have a low polarization resistance, which implies high exchange current density. This happens for high rate constants for the anodic and cathodic reactions and high concentrations of reacting species O and R. With regard to this criterion, the SHE that has i 0 > 10 -3 A/ geometric cm 2 is a particularly good reference electrode (Ref 3). It is a question of judgment how polarizable the reference electrode can be. The answer depends on the precision required and the impedance of the voltmeter used. A high-impedance voltmeter (10 12 ohms) may provide acceptable results with a relatively polarizable electrode. The Liquid Junction Potential. Reference electrodes are usually made of a metal immersed in a well-defined electrolyte. In the case of the CuSO 4 /Cu electrode, the electrolyte is a saturated CuSO 4 aqueous solution; for the SCE, it is a saturated KCl solution. This electrolyte that characterizes the reference electrode must come into contact with the liquid environment of the test electrode to complete the measuring circuit (Fig. 4). There is direct contact between different aqueous media. The difference in chemical composition of the two solutions produces a phenomenon of interdiffusion. In this process, except for a few electrolytes such as KCl, the cations and anions move at different speeds. As an example, in hydrogen chloride (HCl) solution in contact with another medium, the H + ions move faster than the Cl - ions. As a result, a charge separation appears at the limit between the two liquids, the liquid junction. This produces a potential difference called the liquid junction potential, which is included in the measured voltage, V, as expressed in: V = V T - V R + V LJP (Eq 12) where V T is the test potential to be measured, V R is the reference electrode potential, and V LJP is the unknown liquid junction potential. Fig. 4 Schematic of an electrochemical cell with liquid junction potential. P, interface; V, voltmeter In order to determine V T , the liquid junction potential must be eliminated or minimized. The best way, when possible, is to design a reference electrode using an electrolyte identical to the solution in which the test electrode is immersed (Fig. 4). However, in most cases this is not possible, and the best approach is to minimize the liquid junction potential by using a reference electrolyte with a chemical composition as close as possible to the corrosion environment. The use of a solution of KCl (such as in the calomel electrode) offers a partial answer. The diffusion rates of potassium (K + ) and chloride (Cl - ) ions are similar. In contact with another electrolyte, a KCl solution does not produce much charge separation and, consequently, no significant liquid junction potential. The ions present in the other solution, however, also diffuse, and they may do so at different rates, thus producing some separation of charge at the interface (P in Fig. 4). The remaining liquid junction potential, after minimization, constitutes an error that is frequently accepted in electrode potential measurements, especially when compared with results determined under similar experimental conditions. While liquid junction potentials must be minimized as much as possible, there is no general solution for this; each individual case must be well thought out. References cited in this section 1. D.J.G. Ives and J. Janz, Reference Electrodes, Academic Press, 1961 2. C.H. Hamann, A.H. Hamnett, W. Vielstich, Electrochemistry, Wiley-VCH, Weinheim, 1998 3. M.G. Fontana, Corrosion Engineering, 2nd ed., McGraw-Hill, 1978, p 12 E. Protopopoff and P. Marcus, Potential Measurements with Reference Electrodes, Corrosion: Fundamentals, Testing, and Protection, Vol 13A, ASM Handbook, ASM International, 2003, p 13–-16 Potential Measurements with Reference Electrodes E. Protopopoff, Laboratoire de Physico-Chimie des Surfaces, CNRS, and P. Marcus, Ecole Nationale Supérieure de Chimie de Paris, Université Pierre et Marie Curie Operating Conditions for Reference Electrodes When a reference electrode has been selected for a particular application, its proper use requires caution and specific measurement conditions. When measuring the potential of a polarized test electrode versus a reference electrode, it is important not to polarize or damage the latter by applying a significant current density. Also, the ohmic (IR) drop must be minimized. Very Low Current Density. It is important to use a reference electrode that operates at its known open-circuit potential and thus avoid applying any significant overpotential to it. This is achieved by using a high- impedance voltmeter that has a negligible input current and, for test electrode polarization measurements, by using an auxiliary electrode in a three-electrode system (Fig. 1). The value tolerated for the maximum overpotential on the reference electrode, at the condition that it stays under the limit over which the electrode suffers irreversible damages (like the calomel electrode), is a matter of judgment that depends on the accepted magnitude of error in the particular case under investigation. The use of an electrometer or a high-impedance voltmeter (10 12 ohms) fulfills the usual requirements. When a lower impedance instrument is used, an unacceptable overpotential could result if the electrode is too polarizable. The IR Drop and Its Mitigation. The IR drop is an ohmic voltage that results from electric current flow in ionic solutions. Electrolytes have an ohmic resistance; when a current passes through them, an IR voltage can be observed between two distinct points. When the reference electrode is immersed at some distance from a working or test electrode, it is in the electric field somewhere along the current path. An electrolyte resistance exists along the path between the test and the reference electrodes. As current flows through that path, an IR voltage appears in the potential measurement according to: V = V T - V R + IR (Eq 13) where V T is the test potential to be measured, V R is the reference electrode potential, and IR is the ohmic drop. In this case, the liquid junction potential has been neglected. The IR drop constitutes a second unknown value in a single equation. It must be eliminated or minimized. The Luggin capillary is a tube, usually made of glass, that has been narrowed by elongation at one end. The narrow end is placed as close as possible to the test electrode surface (Fig. 1), and the other end of the tube goes to the reference electrode compartment. The Luggin capillary is filled with cell electrolyte, which provides an electric link between the reference and the test electrode. The use of a high-impedance voltmeter prevents significant current flow into the reference electrode and into the capillary tube between the test electrode and the reference electrode compartment (Fig. 1). This absence of current eliminates the IR drop, and the measurement of V T is then possible. A residual IR drop may, however, exist between the tip of the Luggin capillary and the test electrode. This is usually negligible, however, especially in high-conductivity media. The remote electrode technique can be used only for measurement in an electrolyte with very low resistivity, usually in the laboratory. It is applicable, for example, in a molten salt solution, in which the ohmic resistance R is very small. In such a case, the reference electrode can be placed a few centimeters away from the test electrode, because the IR drop remains negligible. In other electrolytes (for example, in measurements in soils), the ohmic resistance is rather large, and the IR drop cannot be eliminated in this manner. The Current Interruption Technique. In this case, when the current is flowing, the IR drop is included in the measurement. A recording of the potential is shown in Fig. 5. At time t 1 , the current is interrupted so that I = 0 and IR = 0. Fig. 5 The potential decay at current interruption. IR is the potential drop due to the electrolyte ohmic resistance. At the moment of the interruption, however, the electrode is still polarized, as can be seen at point P in Fig. 5. The progressive capacitance discharge and depolarization of the test electrode take some time. The potential measured at the instant of interruption then represents the test electrode potential corrected for the IR drop. Precise measurements of this potential are obtained with an oscilloscope. Potential Conversion Between Reference Electrodes. Due to the number of different reference electrodes used, each potential measurement must be accompanied by a clear statement of the reference used. It is often needed to express electrode potentials versus a particular reference, regardless of the actual reference used in the measurement. The procedure is illustrated in the following example. The electrode potential of a buried steel pipe is measured with respect to a CuSO 4 /Cu electrode, and the value is -650 mV for a pH 4 environment. If that value is mistakenly placed in the iron E-pH (Pourbaix) diagram (Fig. 1 in the article “Potential versus pH (Pourbaix) Diagrams” in this Volume), it could be concluded that corrosion will not occur. This conclusion, however, would be incorrect, because the E-pH (Pourbaix) diagrams are always computed with respect to the SHE. It is then necessary to express the measured electrode potential with respect to the SHE before consulting the E-pH (Pourbaix) diagram. The CuSO 4 /Cu electrode potential is +310 mV versus SHE, so this value must be added to the measured potential: E SHE = -650 + 310 = -340 mV. The principle of this conversion is illustrated in the electrode potential conversion diagram of Fig. 6. Fig. 6 Diagram of potential conversion between reference electrodes. SHE, standard hydrogen electrode; CuSO 4 , copper-copper sulfate electrode. SHE, standard hydrogen electrode The value of -340 mV placed in the E-pH Pourbaix diagram at a pH 4 clearly lies in the corrosion region for iron. It would then be definitely necessary to consider the cost benefit of a protection system for the steel pipe. E. Protopopoff and P. Marcus, Potential Measurements with Reference Electrodes, Corrosion: Fundamentals, Testing, and Protection, Vol 13A, ASM Handbook, ASM International, 2003, p 13–-16 Potential Measurements with Reference Electrodes E. Protopopoff, Laboratoire de Physico-Chimie des Surfaces, CNRS, and P. Marcus, Ecole Nationale Supérieure de Chimie de Paris, Université Pierre et Marie Curie Acknowledgment Portions of this article have been adapted from D.L Piron, Potential Measurements with Reference Electrodes, Corrosion, Vol 13, Metals Handbook, 9th ed., ASM International, 1987, p 21–24. E. Protopopoff and P. Marcus, Potential Measurements with Reference Electrodes, Corrosion: Fundamentals, Testing, and Protection, Vol 13A, ASM Handbook, ASM International, 2003, p 13–-16 Potential Measurements with Reference Electrodes E. Protopopoff, Laboratoire de Physico-Chimie des Surfaces, CNRS, and P. Marcus, Ecole Nationale Supérieure de Chimie de Paris, Université Pierre et Marie Curie References 1. D.J.G. Ives and J. Janz, Reference Electrodes, Academic Press, 1961 2. C.H. Hamann, A.H. Hamnett, W. Vielstich, Electrochemistry, Wiley-VCH, Weinheim, 1998 3. M.G. Fontana, Corrosion Engineering, 2nd ed., McGraw-Hill, 1978, p 12 4. L. Pauling, General Chemistry, W.H. Freeman, 1964, p 338–360 E. Protopopoff and P. Marcus, Potential versus pH (Pourbaix) Diagrams, Corrosion: Fundamentals, Testing, and Protection, Vol 13A, ASM Handbook, ASM International, 2003, p 17–-30 Potential versus pH (Pourbaix) Diagrams E. Protopopoff and P. Marcus, CNRS, Ecole Nationale Supérieure de Chimie de Paris, Université Pierre et Marie Curie Introduction THE PRINCIPLE OF POTENTIAL-pH DIAGRAMS was established in the 1940s in Belgium by Marcel Pourbaix (Ref 1, 2, 3, 4). A potential- pH diagram is a graphical representation of the relations, derived from the Nernst equation, between the pH and the equilibrium potentials (E) of the most probable electrochemical reactions occurring in a solution containing a specific element. The standard equilibrium potentials are computed from thermodynamic data (standard chemical potentials, or Gibbs free energies of formation). The equilibrium relations drawn for a given concentration of the element or for a given ratio of activities of two dissolved species of the element give E-pH lines. The representation of the equilibrium pHs for acid-base reactions (independent of the potential) gives vertical lines. All those lines delimit E-pH domains of stability for the various species of the element, metal, ions, oxides, and hydroxides. Potential- pH diagrams synthesize many important types of information that are useful in corrosion and in other fields. They make it possible to discern at a glance the stable species for specific conditions of potential and pH (Ref 1, 2, 3, 4). The principle of E-pH diagrams may be simply understood with the case of iron in water. Corrosion in deaerated water is expressed by the electrochemical reaction Fe → Fe 2+ + 2e - . The equilibrium potential for the Fe 2+ /Fe couple can be calculated using the Nernst equation: (Eq 1) where is the standard potential value for the couple, R is the gas constant, T is the absolute temperature, F is the Faraday constant, and is activity for the ferrous ion in solution. For a given temperature and Fe 2+ concentration (activity ), the equilibrium potential is constant and is represented as a horizontal line in a E-pH diagram (Fig. 1). This line indicates the potential at which Fe and Fe 2+ at a given concentration are in equilibrium and can coexist with no net tendency for one to transform into the other. At potentials above the line, iron metal is not stable and tends to dissolve as Fe 2+ , hence the Fe 2+ concentration increases until a new equilibrium is reached; this is a domain of stability for Fe 2+ . At potentials below the equilibrium line, the stability of the metallic iron increases, Fe 2+ tends to be reduced, and thus its concentration decreases; this is the domain of stability for the metal (Fig. 1). . Protopopoff and P. Marcus, Potential Measurements with Reference Electrodes, Corrosion: Fundamentals, Testing, and Protection, Vol 13A, ASM Handbook, ASM. Protopopoff and P. Marcus, Potential Measurements with Reference Electrodes, Corrosion: Fundamentals, Testing, and Protection, Vol 13A, ASM Handbook, ASM