A.J. Bard, L.R. Faulkner - Electrochemical methods - Fundamentals and Applications (794273), страница 3
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Many electrochemicalmethods have been devised. Their application requires an understanding of the fundamental principles of electrode reactions and the electrical properties of electrode-solution interfaces.In this chapter, the terms and concepts employed in describing electrode reactionsare introduced. In addition, before embarking on a detailed consideration of methodsfor studying electrode processes and the rigorous solutions of the mathematical equations that govern them, we will consider approximate treatments of several differenttypes of electrode reactions to illustrate their main features. The concepts and treatments described here will be considered in a more complete and rigorous way in laterchapters.2 • Chapter 1.
Introduction and Overview of Electrode Processes1.1.1Electrochemical Cells and ReactionsIn electrochemical systems, we are concerned with the processes and factors that affectthe transport of charge across the interface between chemical phases, for example, between an electronic conductor (an electrode) and an ionic conductor (an electrolyte).Throughout this book, we will be concerned with the electrode/electrolyte interface andthe events that occur there when an electric potential is applied and current passes.
Chargeis transported through the electrode by the movement of electrons (and holes). Typicalelectrode materials include solid metals (e.g., Pt, Au), liquid metals (Hg, amalgams), carbon (graphite), and semiconductors (indium-tin oxide, Si). In the electrolyte phase,charge is carried by the movement of ions. The most frequently used electrolytes are liq++uid solutions containing ionic species, such as, H , Na , Cl~, in either water or a nonaqueous solvent.
To be useful in an electrochemical cell, the solvent/electrolyte systemmust be of sufficiently low resistance (i.e., sufficiently conductive) for the electrochemical experiment envisioned. Less conventional electrolytes include fused salts (e.g., moltenNaCl-KCl eutectic) and ionically conductive polymers (e.g., Nation, polyethyleneoxide-LiClO4). Solid electrolytes also exist (e.g., sodium j8-alumina, where charge is carried by mobile sodium ions that move between the aluminum oxide sheets).It is natural to think about events at a single interface, but we will find that one cannotdeal experimentally with such an isolated boundary.
Instead, one must study the properties of collections of interfaces called electrochemical cells. These systems are definedmost generally as two electrodes separated by at least one electrolyte phase.In general, a difference in electric potential can be measured between the electrodes inan electrochemical cell. Typically this is done with a high impedance voltmeter. This cellpotential, measured in volts (V), where 1 V = 1 joule/coulomb (J/C), is a measure of theenergy available to drive charge externally between the electrodes. It is a manifestation ofthe collected differences in electric potential between all of the various phases in the cell.We will find in Chapter 2 that the transition in electric potential in crossing from one conducting phase to another usually occurs almost entirely at the interface.
The sharpness ofthe transition implies that a very high electric field exists at the interface, and one can expect it to exert effects on the behavior of charge carriers (electrons or ions) in the interfacial region. Also, the magnitude of the potential difference at an interface affects therelative energies of the carriers in the two phases; hence it controls the direction andthe rate of charge transfer. Thus, the measurement and control of cell potential is one of themost important aspects of experimental electrochemistry.Before we consider how these operations are carried out, it is useful to set up a shorthand notation for expressing the structures of cells. For example, the cell pictured in Figure 1.1.1a is written compactly asZn/Zn 2 + , СГ/AgCl/Ag(l.l.l)In this notation, a slash represents a phase boundary, and a comma separates two components in the same phase.
A double slash, not yet used here, represents a phase boundarywhose potential is regarded as a negligible component of the overall cell potential. Whena gaseous phase is involved, it is written adjacent to its corresponding conducting element. For example, the cell in Figure 1.1.1ft is written schematically asPt/H2/H+, СГ/AgCl/Ag(1.1.2)The overall chemical reaction taking place in a cell is made up of two independenthalf-reactions, which describe the real chemical changes at the two electrodes. Each halfreaction (and, consequently, the chemical composition of the system near the electrodes)1.1 IntroductionPtZn3H2AgСГСГjExcessAgCI(а)ExcessAgCI(Ь)Figure l.l.l Typical electrochemical cells, (a) Zn metal and Ag wire covered with AgCI immersedin a ZnCl2 solution, (b) Pt wire in a stream of H2 and Ag wire covered with AgCI in HC1 solution.responds to the interfacial potential difference at the corresponding electrode.
Most of thetime, one is interested in only one of these reactions, and the electrode at which it occursis called the working (or indicator) electrode. To focus on it, one standardizes the otherhalf of the cell by using an electrode (called a reference electrode) made up of phaseshaving essentially constant composition.The internationally accepted primary reference is the standard hydrogen electrode(SHE), or normal hydrogen electrode (NHE), which has all components at unit activity:Pt/H2(a - l)/H + (a = 1, aqueous)(1.1.3)Potentials are often measured and quoted with respect to reference electrodes other thanthe NHE, which is not very convenient from an experimental standpoint. A common reference is the saturated calomel electrode (SCE), which isHg/Hg2Cl2/KCl (saturated in water)(1.1.4)Its potential is 0.242 V vs.
NHE. Another is the silver-silver chloride electrode,Ag/AgCl/KCl (saturated in water)(1.1.5)with a potential of 0.197 V vs. NHE. It is common to see potentials identified in the literature as "vs. Ag/AgQ" when this electrode is used.Since the reference electrode has a constant makeup, its potential is fixed. Therefore,any changes in the cell are ascribable to the working electrode. We say that we observe orcontrol the potential of the working electrode with respect to the reference, and that isequivalent to observing or controlling the energy of the electrons within the working electrode (1, 2).
By driving the electrode to more negative potentials (e.g., by connecting abattery or power supply to the cell with its negative side attached to the working electrode), the energy of the electrons is raised. They can reach a level high enough to transferinto vacant electronic states on species in the electrolyte. In that case, a flow of electronsfrom electrode to solution (a reduction current) occurs (Figure 1.1.2a). Similarly, the energy of the electrons can be lowered by imposing a more positive potential, and at somepoint electrons on solutes in the electrolyte will find a more favorable energy on the electrode and will transfer there.
Their flow, from solution to electrode, is an oxidation current (Figure 1.1.2b). The critical potentials at which these processes occur are related tothe standard potentials, E°, for the specific chemical substances in the system.4Chapter 1. Introduction and Overview of Electrode ProcessesElectrodeSolutionSolutionElectrodeSolutionVacantMO0PotentialElectrodeEnergy levelof electrons0jOccupiedMOA + e —> A(a)ElectrodeSolutionVacantMO0Energy levelof electronsPotential0lOccupiedMOA - e -^ A+(b)Figure 1.1.2 Representation of (a) reduction and (b) oxidation process of a species, A, insolution. The molecular orbitals (MO) of species A shown are the highest occupied MO and thelowest vacant MO. These correspond in an approximate way to the E°s of the A/A~ and A + /Acouples, respectively.
The illustrated system could represent an aromatic hydrocarbon (e.g.,9,10-diphenylanthracene) in an aprotic solvent (e.g., acetonitrile) at a platinum electrode.Consider a typical electrochemical experiment where a working electrode and a reference electrode are immersed in a solution, and the potential difference between the electrodes is varied by means of an external power supply (Figure 1.1.3). This variation inpotential, £, can produce a current flow in the external circuit, because electrons cross theelectrode/solution interfaces as reactions occur. Recall that the number of electrons thatcross an interface is related stoichiometrically to the extent of the chemical reaction (i.e.,to the amounts of reactant consumed and product generated).
The number of electrons ismeasured in terms of the total charge, Q, passed in the circuit. Charge is expressed inunits of coulombs (C), where 1 С is equivalent to 6.24 X 10 1 8 electrons. The relationshipbetween charge and amount of product formed is given by Faraday's law; that is, the passage of 96,485.4 С causes 1 equivalent of reaction (e.g., consumption of 1 mole of reactant or production of 1 mole of product in a one-electron reaction). The current, /, is therate of flow of coulombs (or electrons), where a current of 1 ampere (A) is equivalent to 1C/s. When one plots the current as a function of the potential, one obtains a current-potential (i vs. E) curve.
Such curves can be quite informative about the nature of the solutionand the electrodes and about the reactions that occur at the interfaces. Much of the remainder of this book deals with how one obtains and interprets such curves.1.1 Introduction5Powersupply-AgPt-AgBr1МНВГFigure 1.1.3 Schematic diagram of theelectrochemical cell Pt/HBr(l M)/AgBr/Ag attachedto power supply and meters for obtaining a currentpotential (i-E) curve.Let us now consider the particular cell in Figure 1.1.3 and discuss in a qualitativeway the current-potential curve that might be obtained with it.
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