Chemistry
Standard Potential
Standard potential, also known as standard electrode potential, is a measure of the tendency of a chemical species to gain or lose electrons. It is defined relative to a standard hydrogen electrode and is used to compare the reactivity of different redox reactions. The more positive the standard potential, the greater the tendency for reduction to occur.
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10 Key excerpts on "Standard Potential"
eBook - PDFElectrochemical 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.
eBook - PDF- Hubert H. Girault(Author)
- 2002(Publication Date)
- EPFL PRESS(Publisher)
1 Electrochemical Potential CHAPTER 1 ELECTROCHEMICAL POTENTIAL 1.1 ELECTROCHEMICAL POTENTIAL OF IONS The chemical potential is the main thermodynamic tool used to treat chemical equilibria. It allows us to predict whether a reaction can happen spontaneously, or to predict the composition of reactants and products at equilibrium. In this book, we shall consider electrochemical reactions that involve charged species, such as electrons and ions. In order to be able to call on the thermochemical methodology, it is convenient to define first of all the notion of electrochemical potential, which will be the essential tool used for characterising the reactions at electrodes as well as the partition equilibria between phases. To do this, let us recall first of all, what a chemical potential is, and in particular the chemical potential of a species in solution. 1.1.1 Chemical potential Thermodynamic definition Let us consider a phase composed of chemical species j. By adding to this phase one mole of a chemical species i whilst keeping the extensive properties of the phase constant, i.e. the properties linked to its dimensions (V, S, n j ), we increase the internal energy U of the phase. In effect, we are adding the kinetic energy E trans , the rotational energy E rot and the vibrational energy E vib if i is a molecule, the interaction energy between the species E int , perhaps the electronic energy E el if we have excited electronic states and the energy linked to the atomic mass of the atoms E mass if we consider radiochemical aspects, such that: U E E E E E E = + + + + + trans rot vib el mass int (1.1) Thus, we define the chemical potential of the species i as being the increase in inter- nal energy due to the addition of this species µ ∂ ∂ i i V S n U n j i = ≠ , , (1.2) In general, the variation in internal energy can be written in the form of a differential:
eBook - PDF- Thomas F. Fuller, John N. Harb(Authors)
- 2018(Publication Date)
- Wiley(Publisher)
Because it is infeasible to measure and list the Gibbs energy for every reaction under every set of conditions, we tabulate the Gibbs energy at an arbitrary standard condition (25 °C, 1 bar) for species in a reference state (more on this later), and then correct this value for the particular conditions of interest. Combining Equation 2.10 with Equation 2.3 yields U U θ cell RT nF ln ∏a s i i (2.11) where U θ cell is the Standard Potential for the full cell (ΔG ° nFU θ ) at 25 °C and is defined at the same standard conditions as ΔG ° . Also, similar to ΔG ° , U θ cell is not an absolute quantity, but a difference or relative value. One way to determine the Standard Potential for the full cell is to calculate it from the standard change in Gibbs energy for the reaction. However, it is often easier to use tabulated values of the Standard Potential for the reactions of interest as described in the next section. 2.4 Standard PotentialS For convenience, values of Standard Potentials have been determined and tabulated for a large number of electro- chemical reactions, a subset of which is found in Appendix A. Although listed as half-cell reactions, the potentials represent the difference between the potential of the reaction of interest and a reference reaction. The universal reference is the standard hydrogen electrode (SHE). Hence, the stan- dard potential for the hydrogen reaction is defined to be zero. Tables of Standard Potentials are typically written as catho- dic (reduction) reactions. Be mindful that for our thermo- dynamic analysis these reactions are assumed to be in equilibrium. Reactions whose Standard Potentials are posi- tive relative to hydrogen naturally act as cathodes when coupled with hydrogen. In contrast, reactions with negative potentials are more anodic than hydrogen. A table such as that found in Appendix A is also a good place to start when you are unsure of how a particular compound might react.
eBook - PDF- Douglas Skoog, Donald West, F. Holler, Stanley Crouch, Douglas Skoog(Authors)
- 2021(Publication Date)
- Cengage Learning EMEA(Publisher)
16C-6 The Standard Electrode Potential, E 0 When looking carefully at Equations 16-11 and 16-12, note that the constant E 0 is the electrode potential whenever the concentration quotient (actually, the activ- ity quotient) has a value of 1. This constant is by definition the standard electrode potential for the half-reaction. Note that the quotient is always equal to 1 when the activities of the reactants and products of a half-reaction are unity. The standard electrode potential for a half-reaction, E 0 , is defined as the electrode potential when all reactants and products of a half- reaction are at unit activity. The Nernst expression in part (5) of Example 16-2 requires an excess of solid AgCl so that the solution is saturated with the compound at all times. ❮ 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. 426 CHAPTER 16 Introduction to Electrochemistry The standard electrode potential is an important physical constant that provides quantitative information regarding the driving force for a half-cell reaction. 2 The im- portant characteristics of these constants are the following: 1. The standard electrode potential is a relative quantity in the sense that it is the potential of an electrochemical cell in which the reference electrode (left-hand electrode) is the standard hydrogen electrode, whose potential has been assigned a value of 0.000 V. 2. The standard electrode potential for a half-reaction refers exclusively to a reduc- tion reaction, that is, it is a relative reduction potential.
eBook - PDF- Roger Minear(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
At 25°C and one atmosphere total pressure Eq. (8) can be written £ n = £ . + o ^ l o g ? M 2 l _ _ . ( 9 ) η a B a c a D Equation (8), the Nernst equation, expresses the thermodynamic relation-ship between the redox potential (E h ) and the solution composition. The equilibrium redox potential for any reaction can be calculated from the Nernst equation (8) provided that the Standard Potential (E°) for the reaction and the activities of the reactants and products are known. The Standard Potentials for a large number of redox reactions are available (e.g., Latimer, 1952; Chariot, 1958; Clark, 1960). It is also possible to calculate the Standard Potential of a redox reaction using Eq. (5) provided that the Gibbs free energies of participating species are available. A vast body of free energy data is readily accessible (e.g., Robie and Waldbaum, 1968; Sillén and Martell, 1964 and 1971; Rossini et al., 1952; Naumov et al., 1971, and Robie et al., 1979). This provides a powerful and convenient tool for predicting the equilibrium redox potential and/or the equilibrium solution composition of a system. The redox potential can be measured electrochemically in some sys-tems. This method has been used in the experimental determination o f £ R for reactions in which electrochemical reversibility has been demon-strated. In general, however, the measured electrode potential differs from that predicted by equilibrium calculations for one or more of the reasons that will be discussed in Section III. The redox potential is commonly expressed in volts. It should be ap-parent from the discussion above, however, that the redox potential is equivalent to the free energy change per mole of electrons associated with a half-cell reaction and can as readily be expressed in terms of a free 2 Redox Measurements 59 energy or an electron activity.
eBook - PDFElectrochemical Methods
Fundamentals and Applications
- Allen J. Bard, Larry R. Faulkner(Authors)
- 2012(Publication Date)
- Wiley(Publisher)
[See also reference (9), and Sections 1.3.4, and 1.4.2(b).] The Standard Potential of a cell or half-reaction is obtained under conditions where all species are in their standard states (10). For solids, like Ag in cell 2.1.32 or reaction 2.1.33, the standard state is the pure crystalline (bulk) metal. It is interesting to consider how many atoms or what particle size is needed to produce “bulk metal” and whether the Standard Potential is a function of particle size when one deals with metal clusters. These questions have been addressed (11–13); and for clusters containing n atoms (where n 20), indeed turns out to be very different from the value for the bulk metal (n 20). Consider, for example, silver clusters, Ag n . For a silver atom (n 1), the value of can be related to E 0 for the bulk metal through a thermodynamic cycle involving the ionization potential of Ag and the hydration energy of Ag and Ag . This process yields Ag (aq) e L Ag 1 (aq) 1.8 V vs. NHE (2.1.34) which is 2.6 V more negative than for bulk Ag. This result implies that it is much easier energetically to remove an electron from a single isolated Ag atom than to remove an electron from Ag atoms within a lattice of other Ag atoms. Experimental work carried out with larger silver clusters shows that as the cluster size increases, moves toward the value for the bulk metal. For example, for n 2 Ag (aq) Ag 1 (aq) e L Ag 2 (aq) 0 V vs. NHE (2.1.35) These differences in Standard Potential can be explained by the greater surface en- ergy of small clusters compared to bulk metal and is consistent with the tendency of small particles to grow into larger ones (e.g., the dimerization of 2Ag 1 into Ag 2 or the Ostwald ripening of colloidal particles to form precipitates). Surface atoms are bonded to fewer neighbors than atoms within a crystal; thus an extra surface free energy is required to cre- ate additional surface area by subdivision of a metal.
eBook - PDF- 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.
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2014(Publication Date)
- Wiley(Publisher)
In addition, by defining the standard reduction potential of the hydrogen electrode, we saw that standard reduction potentials could be determined. Now we will begin investigating some of the uses of standard cell potentials. Predicting Spontaneous Reactions It’s easy to predict the spontaneous reaction when the substances in two half-reactions, at standard state, are mixed together. This is because we know that the half-reaction with the more positive reduction potential always takes place as written (namely, as a reduction), while the other half-reaction is forced to run in reverse (as an oxidation). Table 19.1 is arranged with the largest positive reduction potential at the top, to the smallest (most negative) at the bottom. The activity series shown in Table 5.3 is also arranged in the same way. We can extend the principles developed in Section 5.6 to predict spontaneous reactions based on the position of the half-reactions in Table 19.1. For any pair of half-reactions, the one higher up in the table has the more positive standard reduction potential and occurs as a reduction. The other half-reaction is reversed and occurs as an oxidation. This concept is useful if we just want to know if a reaction is spontaneous without calculating E ° cell , or as a quick check after solving a problem. Additionally, for a spontaneous reaction, the reactants are found on the left side of the higher half-reaction and on the right side of the lower half-reaction. (This is usually, but not always, true of systems that are not at standard state.) ■ Strictly speaking, the E° values only tell us what to expect under standard conditions. However, only when E ° cell is small can changes in concentration change the direction of the spontaneous reaction. Example 19.3 Predicting a Spontaneous Reaction What spontaneous reaction occurs if Cl 2 and Br 2 are added to a solution that contains both Cl - and Br - ? Assume standard conditions.
eBook - PDF- 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.
eBook - PDF- 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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