Electrode Potential
Electrode potential refers to the measure of the tendency of an electrode to gain or lose electrons when it is in contact with an electrolyte. It is a key concept in electrochemistry and is used to understand and predict the direction of electron flow in chemical reactions. The electrode potential is measured in volts and is essential for determining the feasibility of redox reactions.
8 Key excerpts on "Electrode Potential"
eBook - ePub- Allen J. Bard, Roger Parsons, Joseph Jordan(Authors)
- 2017(Publication Date)
- CRC Press(Publisher)
...2 The Single Electrode Potential: Its Significance and Calculation * ROGER PARSONS † Laboratoire d'Elect rochimie Interfaciale du CNRS, Meudon, France In Chapter 1 the relation of standard Electrode Potentials to standard thermodynamic quantities was considered. It was emphasized there that by considering whole cell reactions there is no difficulty in principle in relating these quantities, and also that these relationships provide all that is required for the solution of practical problems. Nevertheless, the question of single Electrode Potentials and their relation to other single ionic properties has existed since the early days of electrochemistry and interest in this question still remains active. This chapter attempts to clarify some of these problems. I. Electrical Potentials in Real Systems Many of the problems in real systems arise because of the attempt to adapt the classical concept of an electrical potential to a system of real condensed phases. In classical electrostatics the potential is defined in terms of the energy of a test charge and the nature of this test charge is unimportant provided that the potential difference between two points in a uniform medium is considered. This occurs in a number of real situations: for example, two points in vacuum or in a dilute gas or, in the most important example, two pieces of metal of the same composition which make up the terminals of a potentiometer or digital voltmeter. It is this last example which is used every time a cell potential difference is measured as described in Chapter 1. However, the classical definition of potential causes difficulty as soon as an attempt is made to consider the potential difference between two points in different media: for example, between a point in vacuum and a point in a condensed phase...
eBook - ePubSupercapacitors
Materials, Systems, and Applications
- Francois Beguin, Elzbieta Frackowiak, Max Lu, Francois Beguin, Elzbieta Frackowiak(Authors)
- 2013(Publication Date)
- Wiley-VCH(Publisher)
...Δ χ is mainly due to the following: The adsorption of water dipoles on the surface of the metal. Water is a dipole because the oxygen is electronegative and attracts electron density from the H–O σ bonds. Usually, more dipoles are oriented one way than the other, giving rise to a surface potential difference (Δ χ). A good model for a metal is that it is a matrix of ions embedded in a gas of electrons. The electron gas spills out a little way from the surface, resulting in a potential difference. Note that: Almost always Δ χ Δ ψ and then Δ φ ≈ Δ ψ. By convention, Δ = metal – solution. Thus, 1.3.1.7 The Electrochemical Potential () This is the total potential energy of a charged chemical species in a phase; that is, where μ is the chemical potential of the species, z its formal charge, and F is Faraday's constant. 1.3.2 The Electrically Charged Interface or Double Layer 1.3.2.1 The Interface Consider two phases in contact; that is, The interface is the region where the properties of each phase are influenced by the presence of the other phase. 1.3.2.2 Ideally Polarized Electrode Consider No significant current flows until the applied voltage is greater than 1.2 V; then, An electrode is said to be ideally polarized when no charge-transfer reactions occur on it; for example, in the range −1.0 to +0.2 V for Hg in NaCl solution. Such electrodes are preferred for the study of the interface because of the absence of the complications of charge transfer. The voltmeter measures the potential difference between one Hg electrode and the reference electrode (R). If we move the reference electrode back and forth between the two Hg electrodes we find that the voltmeter reading is unchanged. This means that all the applied voltage falls at the electrode/solution interfaces; that is, This would not be the case if current flowed...
eBook - ePub- John Kenkel(Author)
- 2020(Publication Date)
- CRC Press(Publisher)
...A rechargeable battery, when it is positioned in the recharging unit, would be an example of such a cell (see Figure 11.1). Electroanalytical techniques utilize both general types of cells. FIGURE 11.1 (a) A battery with its negative and positive poles connected is a galvanic cell. (b) A rechargeable battery, when positioned in its recharging unit, is an example of an electrolytic cell. 11.2 POTENTIOMETRY Electroanalytical techniques which measure or monitor Electrode Potential utilize the galvanic cell concept. Such techniques fall under the general heading of “potentiometry.” Examples include the pH measurement, ion-selective electrode measurement, and potentiometric titrations. In these techniques, a pair of electrodes is immersed, and the potential, or voltage, of one of the electrodes is measured, hence the name potentiometry. To understand how and why these techniques work, a fundamental knowledge of the Nernst Equation is needed. 11.2.1 The Nernst Equation All oxidation and reduction half-reactions have a certain tendency to occur. Relative tendencies are often listed in text and reference books in the form of a table of “standard reduction potentials,” symbolized E° and having the units of volts. If the tendency of a particular species to be reduced is high, then it will have a positive E°. For example, the reduction of fluorine, F 2 (to F −), which is perhaps the species with the strongest of all tendencies to reduce, has an E° of +2.87 V. This represents what may be the upper limit to the E° values. If the tendency of a particular species to reduce is low, and in fact is more likely to be formed when another species is oxidized, then its E° value will be negative. Such is the case, for example, with all ions of alkali metals and alkaline earth metals. The E° value of lithium ion, Li +1 (being reduced to Li metal), is perhaps the most negative of all, −3.05. Thus the E° values for typical chemical species range from about −3 to about +3 V...
- Franklin Bretschneider, Jan R. de Weille(Authors)
- 2018(Publication Date)
- Academic Press(Publisher)
...Nernst established the relation between these two as: E = RT zF ln (c K) where E is the Electrode Potential, T is the absolute temperature, z is the valence, c is the concentration of the metal ion, and K is the mentioned solution pressure. Although the latter is a hypothetical quantity and is even criticized as to its physical meaning, the differences between metals are very real. Some metals, such as silver, gold, and platinum, are very hard to dissolve (i.e., oxidize) and are called precious metals. Others, such as aluminum, iron, and zinc, are oxidized very easily. Thus, all metals can be ordered according to their relative solution pressures. Hydrogen, although not a metal in the strict sense, fits in this electrochemical series perfectly. Its use as a standard was already mentioned. The list below shows electrode (i.e., half-cell) potentials for a number of reactions, all taken at “normal” (one equivalent per liter) concentrations of the salt. These are called “standard Electrode Potentials” and are abbreviated E 0. System E 0 (Volt; 25°C) Zn ++ /Zn − 0.761 Fe ++ /Fe − 0.440 Pb ++ /Pb − 0.126 H + /H 2 0 by definition Table Continued System E 0 (Volt; 25°C) AgCl/Ag + +0.222 Calomel +0.281 Cu ++ /Cu +0.337 Hg ++ /Hg +0.789 Ag + /Ag +0.799 Au +++ /Au +1.50 Note that the standard hydrogen electrode half-cell has a pH of 0. Since in biochemical systems, pH values are about pH = 7 rather that at 0, biochemists often use Electrode Potentials referred to a hydrogen electrode at pH = 7.0. These are indicated as E 0 ′ and can be found by subtracting 406 mV from the values in the table (why?). Standard potentials of other redox systems, such as metabolic systems, can be expressed in the same way and are called standard redox potentials. These are almost always expressed as E 0 ′ values. As a result of Faradaic processes, these electrode half-cells conduct electric current with very little change of the potential...
- (Author)
- 2013(Publication Date)
- Elsevier(Publisher)
...Even at small electric potentials high currents can be measured [ 1 ]. The resulting difference of the Galvani potentials Δ φ αβ of both phases is therefore characteristic for the analyzed system and depends on the chemical composition of phases and the temperature [ 2, 3 ]. When both phases come into contact the electrochemical reactions at the interface starts. For example, the cations of the electrode will be transferred to the electrolyte with parallel negative charge of the electrode. This negative charge of the electrode prevents the generation of new cations and their transfer to the electrolyte. On the side of the electrolyte the cations or anions generate adsorption layers at the interface (see Section 1.6.1.2). This process is running until no more current is transferred through the interface [ 3 ]. At this stage, the electrochemical equilibrium is reached, the double layer is developed and the Galvani potentials Δ φ αβ (or contact potential) becomes a stable value. The electrochemical equilibrium is therefore then reached, when besides the chemical potential equilibrium resulting from reaction or mass transport also the electrical potentials of reacting species are taken into account for the estimation of equilibrium. Therefore the energy equation (1.1.10 – 1.1.12) in the Section 1.1.2 and the chemical potentials μ i of the species i must be extended...
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
...The variation of cell potential with temperature is determined by the expanded form of the Nernst equation: E cell = E 0 cell – RT / n e – F ln Q = E 0 cell – 2.303 × RT / n e – F log Q (7) where R is the gas constant (8.314 J/K• mol), T is the temperature, and F is the faraday constant. The Faraday constant is the magnitude of electric charge carried by one mole of electrons. The amount of charge carried by one electron is 1.602 × 10 − 19 C, so the amount of charge carried by one mole of electrons is: F = 1.602 × 10 – 19 C / e – 6.022 × 10 23 e – / mol F = 96 485 C / mol (8) At 298 K, 2.303 × RT / F = 0.059 and Eq. (9) reduces to Eq. (6). A galvanic cell that has both half cells with the same composition will have a standard cell potential of zero: E 0 cathode = E 0 anode ; E 0 cell = E 0 cathode − E 0 anode = 0. This only holds for standard conditions of 1 M concentrations. The Nernst equation predicts that if the concentrations of the ions are different in the two half cells (Q ≠ 0), there will still be a voltage generated (E cell ≠ 0). E cell = − 0.059 log Q (9) This is due to the driving force of the oxidation-reduction reaction trying to reach equilibrium. The same reaction occurs in both half cells, but in opposite directions. The cell with the lower concentration becomes the anode, increasing the concentration of the ions, and the cell with the higher concentration becomes the cathode, decreasing the concentration of the ions. This type of cell system is called a concentration cell, a form of galvanic cell where both half cells have the same composition with differing concentrations. But, since an order of magnitude concentration difference in the half cells only produces less than 60 mV, these cells are not typically used to generate electrical energy. Instead, concentration cells are more commonly used in chemical methods of analysis...
eBook - ePubCorrosion Engineering
Principles and Solved Problems
- Branko N. Popov(Author)
- 2015(Publication Date)
- Elsevier(Publisher)
...They can evaluate the conditions for formation of barrier films on the metals, but they cannot estimate their effectiveness in protecting the metal in different environments. It should be noted that the Nernst equation is used to estimate the Electrode Potentials and is based on thermodynamic equations, which are not accurate when the concentration of the electroactive species is close to zero. All metals have a limiting critical value of their activities, (a concentration of < 10 − 6 g-ions per liter), below which the Nernst equation does not agree with the experimentally measured Gibbs free-energy. 2.11.1 Regions of electrochemical stability of water Figure 2.13 shows the thermodynamic stability of water at 25 °C and at standard pressure, as a function of the potential and the pH of the electrolyte. Fig. 2.13 Regions of electrochemical stability of water. The regions of electrochemical stability of water are used to predict the properties of a metal in aqueous solutions when the metal’s potential is known at given pH. In Fig. 2.13, the reversible potential of the hydrogen line, line “b” electrode is constructed using the reversible potential for the hydrogen evolution reaction in an acidic solution: 2 H + + 2 e − → H 2 In an alkaline solution, the equivalent reaction is: 2 H 2 O + 2 e − → H 2 + 2 OH − The half-cell Electrode Potential as a function of the pH of the solution is given as: e H + | H 2 = e H + | H 2 o − 0.059 pH The area below line “b” corresponds to the area where the water decomposition with hydrogen evolution reaction occurs. The area in Fig. 2.13 between the oxygen equilibrium potentials (line “a”) and the hydrogen line (line “b”) is where the water is stable. In this region, water may be synthesized from oxygen and hydrogen...
- Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
...An example of the connection between different chapters is seen in the relationships between free energy, cell potential, and equilibrium constants. Relationship of Cell Potential to Gibbs Free Energy For a voltaic cell, the overall cell potential will always be positive and signifies a spontaneous redox reaction. The relationship between the overall cell potential and the free energy is shown in the following equation: n = number of electrons lost or gained F = Faraday’s constant; 96,500 coulombs/mole E ° = standard cell potential The free energy of the reaction can also be used to calculate the equilibrium constant, K, for a reaction. R = gas constant, 8.31 J/mol K T = absolute temperature, K K = equilibrium constant The combination of the two equations can be used to relate the cell potential of a redox reaction to its equilibrium constant. TEST TIP You should make sure that the units of ΔG are in Joules, not kilojoules, when calculating an equilibrium constant. EXAMPLE: For the following equation, calculate the standard cell potential, the free energy, and the equilibrium constant. SOLUTION: ΔG and Nonstandard Conditions Although the data for standard conditions is prevalent, most reactions do not occur under these specific circumstances. It is still important to predict spontaneity, though, so it will be necessary to calculate ΔG rather than ΔG°. Following is the relationship between ΔG and ΔG°: R = 8.31 J/mol K T = absolute temperature, K Q = reaction quotient Δ G ° = Gibbs free energy under standard conditions ΔG can be related to the cell potential at nonstandard conditions, as well. Here, E is calculated using the Nernst equation. Electrolytic Cell In addition to Voltaic cells, there is another type of electrochemical cell called an electrolytic cell. An electrolytic cell is a nonspontaneous redox reaction that requires energy from an external source...
