Electrode Potential

10 Key excerpts on "Electrode Potential"

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  eBook - PDF

  Fundamentals of Analytical Chemistry

  • Douglas Skoog, Donald West, F. Holler, Stanley Crouch, Douglas Skoog(Authors)
  • 2021(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  An Electrode Potential is by definition a reduction potential. An oxidation potential is the potential for the half-reaction written in the opposite way. The sign of an oxidation potential is, therefore, opposite that for a reduction potential, but the magnitude is the same. Copyright 2022 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 424 CHAPTER 16 Introduction to Electrochemistry The sign of an Electrode Potential is determined by the sign of the half-cell in question when it is coupled to a standard hydrogen electrode. When the half-cell of interest exhibits a positive potential versus the SHE (see Figure 16-7), it will behave spontaneously as the cathode when the cell is discharging. When the half-cell of in- terest is negative versus the SHE (see Figure 16-8), it will behave spontaneously as the anode when the cell is discharging. 16C-5 Effect of Concentration on Electrode Potentials: The Nernst Equation An Electrode Potential is a measure of the extent to which the concentrations of the species in a half-cell differ from their equilibrium values. For example, there is a greater tendency for the process Ag 1 1 e 2 m Ag 1 s 2 to occur in a concentrated solution of silver(I) than in a dilute solution of that ion. It follows that the magnitude of the Electrode Potential for this process must also become larger (more positive) as the silver ion concentration of a solution is increased. Now examine the quantitative relationship between concentration and Electrode Potential.

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  Engineering Chemistry

  Fundamentals and Applications

  • Shikha Agarwal(Author)
  • 2016(Publication Date)
  • Cambridge University Press

    (Publisher)

  Cu 2+ (aq) + 2e –  Cu (s) 2H + (aq) + 2e –  H 2 Figure 14.9 (a) Electrode becomes positive with respect to the solution (Reduction) (b) Electrode becomes negative with respect to the solution (Oxidation) It may be noted that the oxidation potential is the reverse of reduction potential. For example, if the reduction potential of Zn is –0.76 volts, then its oxidation potential is +0.76 volts. If the concentration of the ions is 1.0 M, the temperature is 25°C and for cell reaction where gases are involved, the pressure is one atmosphere, then the Electrode Potential is called the standard Electrode Potential denoted by the symbol E°. The reduction potential under the above-mentioned conditions is called the standard reduction potential. The Electrode Potential depends upon • Nature of metal and its ions • Concentration of the ions in the solution and • Temperature According to the present convention, the half cell potentials are always represented as reduction potentials. EMF or cell potential of a cell An electrochemical cell is obtained by coupling two half cells or electrodes. The electrodes in these half cells have different reduction potentials. Therefore, their tendency to accept electrons is different. The electrode with higher value of reduction potential has a greater tendency to gain electrons, and behaves as a cathode. It acquires electrons from the electrode with lower reduction potential, thereby, forcing it to undergo oxidation and behave as an anode. This difference of Electrode Potential between the two electrodes constituting an electrochemical cell is known as the electromotive force (EMF) or cell potential and is the driving force for the cell reaction. The potential difference is expressed in volts. Therefore, the EMF or cell potential arises from the difference in the tendency of the two ions to get reduced. It is expressed as

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  Electrochemical Methods in Soil and Water Research

  • T.R. Yu, G. L. Ji(Authors)
  • 2016(Publication Date)
  • Pergamon

    (Publisher)

  Electrochemical reactions can occur spontaneously in a galvanic cell (Fig. 1.2). To an electrolytic cell, electrons can be introduced by an externally applied voltage. In a cell for the determination of ion activities shown in Fig. 1.3, unless the resistance in the measuring circuit is infinitely large, there is always a small current coming from the input terminals of the mV-meter. Therefore, for practical electrochemical cells the presence of a current and thus the occurrence of an electrode polarization are very frequently encountered. The potential of an electrode is determined by the concentration of charged particles (ions, electrons) at the electrode surface. When a dynamic equilibrium is attained, the concentration of ions or electrons at the electrode surface is constant. Therefore, according to the Nernst equation, the Electrode Potential should also be constant. When an electric current flows through the electrode, some electrochemi-cal reactions should occur at the interface. For example, when a current flows through a copper electrode dipped in a solution containing copper ions with the direction that the current flows from solution to copper electrode, owing to a slower rate of ion transfer from ambient solution to electrode than the consumption rate at the electrode, the concentration of copper ions at the electrode surface would be reduced. As a result, the Electrode Potential would be lower than its equilibrium value. Conversely, if the current is of the opposite direction, the Electrode Potential may be higher than its equilibrium value. In order to characterize the extent of electrode polarization, a term overpotential (η) is usually used. Overpotential is the difference in Electrode Potential before and after electrode polarization: 1i = E T - E E (1-71) where η χ is the overpotential at an electrode current density i, E x the Electrode Potential at that current density, and E c is the equilibrium potential of the electrode.

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  An Introduction to Chemical Metallurgy

  International Series on Materials Science and Technology

  • R. H. Parker, D. W. Hopkins(Authors)
  • 2016(Publication Date)
  • Pergamon

    (Publisher)

  Electrode Potentials calculated on the basis of (5.23) are therefore called reduction potentials. If the electrode re-actions were written down in the reverse direction, the Electrode Potentials would be those of oxidation reactions, and are called oxidation potentials. The sign of E° would be reversed, the logarithm of the reciprocal of the ionic activity would be used, and (5.23) would become E = E°-^na M z+. (5.32) zF The convention chosen here gives the correct sign convention according to the European system and leads to the expres-sion in (5.23) of reduction potentials at electrodes. This system has been adopted by the International Union of Pure and Applied Chemists (I.U.P.A.C.), and is commonly used by electrochemists and corrosion scientists. Unfortunately many American workers and textbooks use the American con-vention, which considers oxidation potentials by (5.23), so that care must be taken before using data from these sources. A fourth type of reaction can occur at an electrode, the reduction of an ionic species which remains in the electrolyte ELECTROCHEMISTRY 195 while the metal of the electrode remains inert. An example of this is a platinum electrode dipping into an aqueous solution of ferrous and ferric ions —two valency states of the same metal, PtlFe 2+ Fe 3+ The reaction at the electrode, known as a redox reaction, is p e 3 + i € - — p e 2+ 1 ^aq ' c *■ w a q ' the ferric ion being reduced to the ferrous state. The redox potential of this electrode is given by the Nernst equation (5.23), £ = E° + ^ l n ^ ^ , (5.33) F 0 F e 2 + which can be generalized as p— po _L ΕΣΛ [activity of reacting species in oxidized state] zF [activity of reacting species in reduced state] (5.34) In steelmaking, the transfer of sulphur from the molten steel to the molten slag can be considered to be a redox reaction [S] + 2e- = (S 2 -), where square brackets indicate a constituent of the metallic phase, round brackets a constituent of the slag phase.

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  Electrochemical Kinetics

  Theoretical Aspects

  • Klaus J. Vetter(Author)
  • 2013(Publication Date)
  • Academic Press

    (Publisher)

  Before the development of specific concepts about the formation of the metal/electrolyte potential difference by means of electrode kinetics, this process was interpreted by Nernstian thermodynamics. According to Nernst, metals have a certain tendency to send their ions into solution as solvated particles. This tendency is expressed as a solution pressure or a solution tension which has a charac-teristic value for each metal. This pressure is opposed by the osmotic pressure of the dissolved solvated metal ions in the electrolyte solution. A potential difference is formed by the difference of the two pressures. If both pressures are identical, an equilibrium exists without poten-tial difference. In the case of Fig. 5a, the solution tension is greater than the osmotic p r e s -sure, so that positive metal ions pass into the electrolyte, leaving behind the negative counter charge (electrons). The created potential FIG . 5a FIG . 5b difference (Fig. 5a) decreases the solution pressure until at further increase of potential difference the decreased solution tension becomes equal to the osmotic pressure. Thus, equilibrium is reached at a specific equilibrium potential. In the case of Fig. 5b, the solution tension is smaller than the osmotic pressure. Here, ions pass from the electrolyte onto the metal with the charge distribution shown in Fig. 5b, until the po-tential difference again causes both pressures to become equal. According to Nernst, the decrease in potential occurs in both cases across the electrical double layer which forms at the surface. The sign and size of the charges are unknown. 19. Concentration Dependence o f the Metal/Ion Potentials The concentration dependence of metal/ion potentials with reference to the standard hydrogen electrode is given by the Nernst equation (1.28), Section 15. A metal/ion electrode must be used in conjunction with the standard hydrogen electrode to form a com-plete cell.

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  Analytical Chemistry

  • Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
  • 2020(Publication Date)
  • Wiley

    (Publisher)

  Chapter 19 ELECTROCHEMICAL CELLS AND Electrode PotentialS “I know of nothing sublime which is not some modification of power.” —Edmund Burke KEY THINGS TO LEARN FROM THIS CHAPTER Voltaic cells Using standard potentials to predict reactions Anodes, cathodes, and cell voltages (key Equation: 19.19) The Nernst equation (key Equation: 19.22) Calculating Electrode Potentials before and after reaction The formal potential An important class of titrations is reduction–oxidation or “redox” titrations, in which an oxidizing agent and a reducing agent react (see Equation 19.1). We define an oxi- dation as a loss of electrons to an oxidizing agent (which itself gets reduced) to give a higher or more positive oxidation state, and we define reduction as a gain of electrons from a reducing agent (which itself gets oxidized) to give a lower or more negative oxidation state. We can gain an understanding of these reactions from a knowledge of electrochemical cells and Electrode Potentials. In this chapter, we discuss electrochem- ical cells, standard Electrode Potentials, the Nernst equation (which describes Electrode Potentials), and limitations of those potentials. Chapter 20 discusses potentiometry, the use of potential measurements for determining concentration, including the glass pH electrode and ion-selective electrodes. In Chapter 21, we describe redox titrations and potentiometric titrations in which potentiometric measurements are used to detect the end point. We review in that chapter the balancing of redox reactions since this is required for volumetric calculations. You may wish to review that material now. Oxidation is a loss of electrons. Reduction is a gain of electrons. “OIL RIG” is a good mnemonic to help remember this. Electrochemical methods, with the exception of the nearly universal use of the potentiometric pH meter, are generally not as widely used as spectrochemical or chro- matographic methods.

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  Inorganic Species, Part 1

  • Roger Minear(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  The electrode function, however, depends upon complete equilibration between the reference solution and the system to be mea-sured. The inherent problems associated with this type of electrode when used for redox measurements in aqueous systems has been discussed by Ben-Yaakov and Kaplan (1973). The ease with which electrode measurements can be made belies the complexity of the electrochemical reactions taking place in the system. The fact that redox measurements made in strongly oxidizing environ-ments are usually large and positive, whereas redox measurements made in strongly reducing environments are usually negative, has lulled some investigators into the dangerous assumption that the measured potentials represent reversible Nernst potentials. Under certain well-controlled con-ditions this may be the case; in general, however, quantitative interpreta-tion of the measured potentials is not possible (Bricker, 1965; Berner, 2 Redox Measurements 61 1963; Stumm, 1967; Morris and Stumm, 1967; Stumm and Morgan, 1981; Whitfield, 1974). A brief examination of the origin of the Electrode Potential will illustrate some of the problems associated with making electrochemically reversible electrode measurements. Consider an electrode in a system containing a redox couple: When the rate of reduction of A 0 to A R is equal to the rate of oxidation of A R to A 0 the reaction is at equilibrium, the flow of current associated with the reduction exactly counterbalances the flow of current associated with the oxidation, and the potential of the reaction is fixed. Although the net current flow is zero, the current flow associated with either the reduction or the oxidation is not. This unidirectional current flow is called the exchange current. If it behaves as an ideal redox elec-trode, the electrode immersed in the system acts as an electron exchanger with both the oxidized and the reduced species participating in the reac-tion, and it assumes the value of the equilibrium potential.

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  Analytical Chemistry

  • Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
  • 2013(Publication Date)
  • Wiley

    (Publisher)

  For example, a mixture of Fe 2+ and Fe 3+ can be analyzed electrochemically. Potentiometric methods are unique in that they measure the activity of a chemical species rather than concentration, and can detect the free form of a metal ion, some of which may be tied up in complexation. In the next four chapters, we will explore these different techniques and their applicability. 383 384 CHAPTER 12 ELECTROCHEMICAL CELLS AND Electrode PotentialS Before beginning our discussions, it is helpful to describe some fundamental electrochemical terms, some of which will be discussed later on: Potential (volts) originates from separation of charge Ohm’s law: E = iR, where E potential in volts, i is current in amperes, and R is resistance in ohms P = Ei, where P is power in volt amperes The first law of thermodynamics says that there is a conservation of energy in a chemical reaction. Work is done by the system or on the system to the extent of energy evolved or absorbed, work being equal to qE, where q is charge, in coulombs transferred across a potential difference E. In the case of electrochemical cells, free energy change, G, is related to the electrical work done in the cell. If E cell is the electromotive force (emf) of the cell and n moles of electrons are involved in the reaction, the electrical work done will be G = −nFE cell , where F is the Faraday constant, 96,487 coulombs/equivalent. If reactants and products are in their standard states, then G ◦ = −nFE ◦ , where E ◦ is the standard cell potential. While our emphasis in these chapters is on analytical applications of redox reactions and electrochemistry, such reactions are important in biochemistry, energy conversion, batteries, environmental chemistry, and other aspects of our lives, and your knowledge of basic redox chemistry and electrochemistry will be helpful in your understanding of these processes.

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  Fundamentals of Electrochemical Science

  • Keith Oldham, Jan Myland(Authors)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  In fact, there is no straightforward way of measuring (or even defining!) a single Electrode Potential. Nevertheless, the concept of an Electrode Potential is such a useful one that electrochemists have invented a way to get around the fact that single Electrode Potentials cannot be measured. What is done is to select one electrode as a standard and define its potential as zero. The universally accepted standard is the standard hydrogen electrode (SHE), at which the electrode process is 4:1:2 H + (aq,a = l) ±e ' ViH 2 (g, 10 5 Pa) Because this electrode is defined to have zero potential, the voltmeter in the diagram below is said to measure the Electrode Potential £ w o r k of the working electrode. Of course, it really does nothing of the kind: in reality the voltmeter measures a cell voltage, as always. But by convention we all agree to speak of this potential difference as the potential of the WE. red ^ black ΔΕ = £ W O r k 4:2 Reference electrodes In effect, we split the measured voltage AE across the cell into the difference of £ w o r k and £ r e f (we here ignore junction potential differences) as in equation 4:1:1, by imagining an extra electrode — a standard hydrogen electrode - to be dipped into the solution as shown below. Of course, we don't actually need to use a SHE: provided that we know £ r e f and measure AE, we can always calculate £ w o r k as AE + £ref . 4:2 Reference electrodes The reference electrode is the one that we are not interested in. We want to be able to incorporate it into our cell and then forget about it. An ideal reference electrode would be one which maintains a constant potential £ r e f , whether we treat it as an anode or a cathode and irrespective of the current (if any) that we pass through it. To come close to this ideal requires, among other things, that there be an abundant supply of all the species involved in the reference electrode reaction and that the activities of all these species be constant.

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  Electrochemical Methods

  Fundamentals and Applications

  • Allen J. Bard, Larry R. Faulkner, Henry S. White(Authors)
  • 2022(Publication Date)
  • Wiley

    (Publisher)

  It is important to grasp that electrostatic potential (not emf ) is the measurable experimental variable. When a half-reaction is chemically reversible, the potential of its electrode will usually remain in the same neighborhood and retain the same sign, whether the reaction proceeds as an oxidation or a reduction (9) [Sections 1.1.7( f ) and 1.3.2( b )]. 2.1.5 Standard States and Activity Standard thermodynamic quantities (e.g., standard free energies of formation for various substances) correspond to situations in which all species of interest, including all participants in any reaction, are in their standard states . Likewise, the standard potential of a cell or half-reaction is defined for conditions where all species are in their standard states (10). This provision is necessary because many thermodynamic quantities, such as the free energy of a solute, depend on concentration. To provide meaningful tabulations, one must specify the concentrations. It is helpful to adopt a standard practice—that is, to define standard states systematically. Standard thermodynamic quantities are marked with a superscript zero, e.g., Δ G 0 (the standard free energy change for a reaction) or E 0 (the standard potential for an electrode). The definitions of standard state are straightforward, but idealized: 5 • For a pure solid or liquid, the standard state is the substance under the standard pressure of 10 5 Pa (1 bar). For solids, it is necessary to specify the allotrope or crystalline form. • For a pure gas, the standard state is the gas under standard pressure, but behaving as an ideal gas. • For a solute, the standard state is a solution of standard concentration under standard pressure, but behaving as though each solute molecule is infinitely diluted.

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