Chemistry
Electrochemical Series
The electrochemical series is a list of metals and nonmetals arranged in order of their standard electrode potentials. This series helps predict the direction of redox reactions and the relative reactivity of different substances. It is a valuable tool in understanding and predicting the behavior of elements in electrochemical processes such as corrosion and electrolysis.
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11 Key excerpts on "Electrochemical Series"
eBook - PDFEngineering Chemistry
Fundamentals and Applications
- Shikha Agarwal(Author)
- 2016(Publication Date)
- Cambridge University Press(Publisher)
Other reference electrodes commonly in use are discussed in Section 14.13. 14.10 Electrochemical Series The arrangement of elements in the order of increasing reduction potential values is called Electrochemical Series. It is also called the activity series. The Electrochemical Series of some typical electrodes with their electrode potentials is tabulated in Table 14.3. Table 14.3 Standard reduction electrode potential at 25°C Electrode Electrode reaction (reduction) E 0 volts Li Li + (aq)+ e – → Li(s) –3.05 K K + (aq) + e – → K(s) –2.93 Ba Ba 2+ (aq) +2e – → Ba(s) –2.90 Ca Ca 2+ (aq) +2e – → Ca(s) –2.87 Na Na + (aq) +e – → Na(s) –2.71 Mg Mg 2+ (aq)+ 2e – → Mg(s) –2.37 Al Al 3+ (aq) + 3e – → Al(s) –1.66 Zn Zn 2+ (aq) +2e – → Zn(s) –0.76 Electrochemistry 765 Cr Cr 3+ (aq) +3e – → Cr (s) –0.74 Fe Fe 2+ (aq) +2e – → Fe(s) –0.44 Cd Cd 2+ (aq)+ 2e – → Cd(s) –0.40 Co Co 2+ (aq)+2e – → Co(s) –0.28 Ni Ni 2+ (aq) +2e – → Ni(s) –0.25 Sn Sn 2+ (aq) +2e – → Sn(s) –0.14 Pb Pb 2+ (aq) +2e – → Pb(s) –0.13 Pt 2H + (aq) +2e – → H 2 (g) 0.00 Cu Cu 2+ (aq) +2e – → Cu(s) +0.34 I 2 I 2 (s) + 2e – → 2I – (aq) +0.54 Fe Fe 3+ (aq) +e – → Fe 2+ (aq) +0.77 Hg Hg 2 2+ (aq) +2e – → Hg(l) +0.79 Ag Ag + (aq) +e – → Ag(s) +0.80 Br 2 Br 2 (l) +2e – → 2Br – (aq) +1.08 Cl Cl 2 (g) +2e – → 2Cl – (aq) +1.36 Au Au 3+ (aq)+3e – → Au(s) +1.42 Mn MnO 4 – (aq)+8H 3 O + (aq)+5e – → Mn 2+ (aq) +12H 2 O(l) +1.51 F 2 F 2 (g) +2e – → 2F – (aq) +2.87 Applications of the Electrochemical Series 1. Relative strength of the oxidizing and reducing agents The Electrochemical Series helps in predicting the oxidizing and reducing ability of a substance. In the Electrochemical Series, the elements are arranged in the order of their increasing reduction potential, hence, the elements situated at the bottom of the series have a greater tendency to get reduced—they are good oxidizing agents. On the other hand, elements at the top of the table have low reduction potential and have a lesser tendency of getting reduced.
eBook - PDFEngineering Chemistry
Fundamentals and Applications
- Shikha Agarwal(Author)
- 2019(Publication Date)
- Cambridge University Press(Publisher)
Applications of the Electrochemical Series 1. Relative strength of the oxidizing and reducing agents The Electrochemical Series helps in predicting the oxidizing and reducing ability of a substance. In the Electrochemical Series, the elements are arranged in the order of their increasing reduction potential; hence, the elements situated at the bottom of the series have a greater tendency to get reduced—they are good oxidizing agents. On the other hand, elements at the top of the table have low reduction potential and have a lesser tendency of getting reduced. Consequently, they may get oxidized and are good reducing agents. Electrochemistry 915 2. Calculation of cell potential or EMF of the cell Cell potential of the cell is the difference of the reduction potential of the cathode and the anode. E ° cell = E ° (cathode) – E ° (anode) 3. Predicting spontaneity or feasibility of a reaction For a cell reaction to be spontaneous, the EMF of the cell as calculated above should be positive. If E ° cell is negative, then the cell reaction will not be feasible. 4. To predict whether a metal will react with acids to give H 2 gas Those metals whose reduction potential is less than the reduction potential of hydrogen can liberate H 2 gas. Such metals are placed above hydrogen in the Electrochemical Series. Hydrogen having higher reduction potential has a greater tendency to accept electrons and get reduced. M → M + (aq) + e – (Monovalent) H + (aq) + e – → H or ½ H 2 Metals placed below hydrogen in the Electrochemical Series have a higher value of reduction potential than hydrogen. Consequently, such metals do not liberate hydrogen on reacting with acids. 5. Replacement tendency The metal on the top of the Electrochemical Series has the tendency to replace the one below it from its solution. This is because the metal placed at the bottom has a greater reduction potential and hence a greater tendency to accept electrons.
eBook - PDF- Keith Oldham, Jan Myland(Authors)
- 2012(Publication Date)
- Academic Press(Publisher)
Thermodynamics implies that the Nernst law always holds when no current passes through an electrode, but in the form RT c n 4:4:26 Ε = E 0 ' -— In — nF r * c o the Nernst law may hold, not only at the null potential, but also during the passage of current. The conditions under which this is the case are discussed in Section 5:8 and Chapter 6. 4:5 Pourbaix diagrams A concept with a distinguished history is the Electrochemical Series of the chemical elements. In this the elements are arranged in order of the standard electrode potentials of their redox couples with an aqueous ion. The figure to the right is a diagram portraying the Electrochemical Series of eleven elements. Notice that the more electronegative elements have the most positive E° values and therefore occur high in the series, whereas the electropositive metals are found towards the bottom. The farther apart two elements are in the Electrochemical Series, the more energetic is the reaction between the oxidized member of the higher element and the reduced member of the lower element. As examples, the displacement reaction of silver ion in aqueous solution with magnesium powder E°/W 2 . 8 9 -1.36-0.80-0.34-0.00--0.40--0.76--1.18--1.68--2.36--3.05-4:5:1 2Ag + (aq) + Mg(s) 2Ag(s) + Mg 2 + (aq) occurs violently and to virtual completion, whereas the reaction 4:5 Pourbaix diagrams 127 4:5:2 Cd 2 + (aq) + Fe(s) > Cd(s) + Fe 2 + (aq) between members of couples that are adjacent in the Electrochemical Series*, is mild and incomplete. In this diagram of the Electrochemical Series the formulas of the two members of each redox couple have been placed on opposite sides of a horizontal line. Taking the Ag + (aq)/Ag(s) couple as an example, this arrangement is intended to suggest 1.2 + 0.8 E/W 0.4 4-A S ( S ) predominant 0.0+ the predominance diagram shown here. The significance of this diagram is that if a working electrode is in contact with both .
eBook - PDF- Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
INTRODUCTION CHAPTER 16 Electrochemistry 16.1 Review of Redox Chemistry 16.2 Galvanic Cells 16.3 Electrode and Cell Potentials 16.4 Potential, Free Energy, and Equilibrium 16.5 Batteries and Fuel Cells 16.6 Corrosion 16.7 Electrolysis Another chapter in this text introduced the chemistry of reduction-oxidation (redox) reactions. This important reaction class is defined by changes in oxidation states for one or more reactant elements, and it includes a subset of reactions involving the transfer of electrons between reactant species. Around the turn of the nineteenth century, chemists began exploring ways these electrons could be transferred indirectly via an external circuit rather than directly via intimate contact of redox reactants. In the two centuries since, the field of electrochemistry has evolved to yield significant insights on the fundamental aspects of redox chemistry as well as a wealth of technologies ranging from industrial-scale metallurgical processes to robust, rechargeable batteries for electric vehicles ( Figure 16.1). In this chapter, the essential concepts of electrochemistry will be addressed. Figure 16.1 Electric vehicles are powered by batteries, devices that harness the energy of spontaneous redox reactions. (credit: modification of work by Robert Couse-Baker) CHAPTER OUTLINE 16.1 Review of Redox Chemistry LEARNING OBJECTIVES By the end of this section, you will be able to: • Describe defining traits of redox chemistry • Identify the oxidant and reductant of a redox reaction • Balance chemical equations for redox reactions using the half-reaction method Since reactions involving electron transfer are essential to the topic of electrochemistry, a brief review of redox chemistry is provided here that summarizes and extends the content of an earlier text chapter (see chapter on reaction stoichiometry). Readers wishing additional review are referred to the text chapter on reaction stoichiometry.
eBook - PDF- Ageetha Vanamudan(Author)
- 2023(Publication Date)
- Delve Publishing(Publisher)
258 14.18 Metallurgy ................................................................................... 259 14.19 Production of Chemicals.............................................................. 259 14.20 Investigations Into the Natural World ........................................... 260 Introduction to Molecular Science 238 In one of two ways, chemical reactions can absorb or release electricity as energy. Electrochemistry, a branch of chemistry that investigates the process, is concerned with the conversion of chemical and electrical energy. Electrochemistry has a wide range of practical applications in everyday life. Batteries, which are commonly used in flashlights, calculators, and vehicles, generate energy through chemical reactions (Shin et al., 2019). Electricity is used to apply metals such as gold and chromium to objects. Electrochemistry plays a significant role in the propagation of nerve impulses in biological systems. Redox chemistry, or the movement of electrons, underpins all electrochemical reactions. Within an electrochemical cell, electricity and chemical energy may be converted back and forth. The following are the three elements of an electrochemical reaction. Only in the presence of a solution that allows for redox reactions can they occur. Conducting these reactions in water facilitates electron and ion mobility. Electron transfer necessitates the use of a conductor. Electrons may go via this conductor, which is often a wire. The salt bridge must be capable of transferring ions through it as part of its function. 14.1 THE ELECTROCHEMISTRY TECHNIQUE The characteristics of the negatively charged electron impact the interactions between matter and electric current. Atoms, groups of molecules, and the like have a strong attraction to positively charged matter particles like protons. This affinity is similar to the chemical attraction between atoms.
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
INTRODUCTION CHAPTER 17 Electrochemistry 17.1 Review of Redox Chemistry 17.2 Galvanic Cells 17.3 Electrode and Cell Potentials 17.4 Potential, Free Energy, and Equilibrium 17.5 Batteries and Fuel Cells 17.6 Corrosion 17.7 Electrolysis Another chapter in this text introduced the chemistry of reduction-oxidation (redox) reactions. This important reaction class is defined by changes in oxidation states for one or more reactant elements, and it includes a subset of reactions involving the transfer of electrons between reactant species. Around the turn of the nineteenth century, chemists began exploring ways these electrons could be transferred indirectly via an external circuit rather than directly via intimate contact of redox reactants. In the two centuries since, the field of electrochemistry has evolved to yield significant insights on the fundamental aspects of redox chemistry as well as a wealth of technologies ranging from industrial-scale metallurgical processes to robust, rechargeable batteries for electric vehicles ( Figure 17.1). In this chapter, the essential concepts of electrochemistry will be addressed. Figure 17.1 Electric vehicles are powered by batteries, devices that harness the energy of spontaneous redox reactions. (credit: modification of work by Robert Couse-Baker) CHAPTER OUTLINE 17.1 Review of Redox Chemistry LEARNING OBJECTIVES By the end of this section, you will be able to: • Describe defining traits of redox chemistry • Identify the oxidant and reductant of a redox reaction • Balance chemical equations for redox reactions using the half-reaction method Since reactions involving electron transfer are essential to the topic of electrochemistry, a brief review of redox chemistry is provided here that summarizes and extends the content of an earlier text chapter (see chapter on reaction stoichiometry). Readers wishing additional review are referred to the text chapter on reaction stoichiometry.
eBook - PDFChemistry
Principles and Reactions
- William Masterton, Cecile Hurley(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
430 17 ▼ Electrochemistry Most “silver” tableware like the ones in the painting are not made of pure silver but rather are silver-plated using an electrolytic cell. If by fire, Of sooty coal th’ empiric Alchymist Can turn, or holds it possible to turn Metals of drossiest Ore to Perfect Gold. —JOHN MILTON Paradise Lost (Book V, Lines 439–442) E lectrochemistry is the study of the interconversion of electrical and chemical energy. This conversion takes place in an electrochemical cell that may be a(n) ■ ■ voltaic (galvanic) cell (Section 17-2), in which a spontaneous reaction generates electrical energy. ■ ■ electrolytic cell (Section 17-6), in which electrical energy is used to bring about a nonspontaneous reaction. All of the reactions considered in this chapter are of the oxidation-reduction type. You will recall from Chapter 4 that such a reaction can be split into two half-reactions. In one half-reaction, referred to as reduction, electrons are consumed; in the other, called oxidation, electrons are produced. There can be no net change in the number of electrons; the number of electrons consumed in reduction must be exactly equal to the number produced in the oxidation half-reaction. These two half-reactions combine to give a balanced redox reaction (Section 17-1). In an electrochemical cell, these two half-reactions occur at two different elec-trodes, which most often consist of metal plates or wires. Reduction occurs at the cathode; ▲ a typical half-reaction might be cathode: Cu 2 1 ( aq ) 1 2 e 2 9: Cu( s ) Oxidation takes place at the anode, ▲ where a species such as zinc metal produces electrons: anode: Zn( s ) 9: Zn 2 1 ( aq ) 1 2 e 2 It is always true that in an electrochemical cell, anions move to the anode; cations move to the cathode . One of the most important characteristics of a cell is its voltage, which is a mea-sure of reaction spontaneity. Cell voltages depend on the nature of the half-reactions Cathode 5 reduction; anode 5 oxi-dation.
eBook - PDF- Arthur Adamson(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
CHAPTER THIRTEEN ELECTROCHEMICAL CELLS The preceding chapter dealt primarily with the physical chemistry of electrolyte solutions; we now concern ourselves with the overall chemical process that occurs when electricity is passed through a conducting solution. The emphasis will be on the work associated with this overall change, as measured by the reversible cell potential. Since reversible work at constant temperature and pressure corre-sponds to a free energy change, we will thus be able to bring the emf of cells into the general scheme of thermodynamics. The chapter concludes with a discussion of irreversible electrode processes, that is, with the physical chemistry of the approach of ions to, and their reaction àt, the surface of an electrode. 13-1 Definitions and Fundamental Relationships A. Cell Conventions An electrochemical cell has, as essential features, a current-carrying solution and two electrodes at which oxidation and reduction processes occur, respectively, as current flows. Figure 13-1 gives a schematic example of a fairly typical cell for this chapter; we have hydrogen and silver-silver chloride electrodes dipping into an aqueous solution of HCl. The hydrogen electrode, incidentally, typically consists of a platinized platinum metal surface arranged so that hydrogen gas bubbles past as it dips partly into the solution, the object being to provide the most intimate possible gas-solution-metal contact. Platinized platinum is merely platinum metal on which additional, very finely divided platinum has been deposited electrolytically; the result is a high area, catalytically active surface. A silver-silver chloride electrode consists of silver on which a fine-grained, adherent deposit of silver chloride has been placed, again electrolytically.
eBook - PDF- Arther Adamson(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
CHAPTER THIRTEEN ELECTROCHEMICAL CELLS The preceding chapter dealt primarily with the physical chemistry of electrolyte solutions; we now concern ourselves with the overall chemical process that occurs when electricity is passed through a conducting solution. The emphasis will be on the work associated with this overall change, as measured by the reversible cell potential. Since reversible work at constant temperature and pressure corre-sponds to a free energy change, we will thus be able to bring the emf of cells into the general scheme of thermodynamics. The chapter concludes with a discussion of irreversible electrode processes, that is, with the physical chemistry of the approach of ions to, and their reaction at, the surface of an electrode. 13-1 Definitions and Fundamental Relationships A. Cell Conventions An electrochemical cell has, as essential features, a current-carrying solution and two electrodes at which oxidation and reduction processes occur, respectively, as current flows. Figure 13-1 gives a schematic example of a fairly typical cell for this chapter; we have hydrogen and silver-silver chloride electrodes dipping into an aqueous solution of HCl. The hydrogen electrode, incidentally, typically consists of a platinized platinum metal surface arranged so that hydrogen gas bubbles past as it dips partly into the solution, the object being to provide the most intimate possible gas-solution-metal contact. Platinized platinum is merely platinum metal on which additional, very finely divided platinum has been deposited electrolytically; the result is a high area, catalytically active surface. A silver-silver chloride electrode consists of silver on which a fine-grained, adherent deposit of silver chloride has been placed, again electrolytically.
eBook - PDF- Young, William Vining, Roberta Day, Beatrice Botch(Authors)
- 2017(Publication Date)
- Cengage Learning EMEA(Publisher)
AlbertSmirnov/iStockphoto.com 21 Electrochemistry Unit Outline 21.1 Oxidation–Reduction Reactions and Electrochemical Cells 21.2 Cell Potentials, Free Energy, and Equilibria 21.3 Electrolysis 21.4 Applications of Electrochemistry: Batteries and Corrosion In This Unit… Why aren’t there more electric cars? Why do we still use corded power tools instead of only using battery-powered tools? Why do the batteries in our portable electronic devices run down so quickly? The answer to all these questions lies in our ability to make good batteries. In this unit we explore the chemistry of batteries, where spontaneous reactions take place by the indirect transfer of electrons from one reactant to another. We will also investigate electrolysis, the process where we use external power supplies such as batteries to force nonspontaneous reactions to form products. Copyright 2018 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. WCN 02-300 Unit 21 Electrochemistry 666 21.1 Oxidation–Reduction Reactions and Electrochemical Cells 21.1a Overview of Oxidation–Reduction Reactions In Chemical Reactions and Solution Stoichiometry (Unit 9) we first encountered oxida-tion–reduction reactions in our study of chemical reactions. Electrochemistry is the area of chemistry that studies oxidation–reduction reactions, also called redox reactions, which involve electron transfer between two or more species. Recall that in a redox reaction, ● the species that loses electrons has been oxidized and is the reducing agent in the reaction, and ● the species that gains electrons has been reduced and is the oxidizing agent in the reaction. Oxidizing and reducing agents are identified in a chemical reaction by using the oxida-tion number (or oxidation state) of the species in the reaction. Recall that the oxidation number of an oxidized species increases during the reaction, whereas the oxidation num-ber of a reduced species decreases during the reaction.
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2013(Publication Date)
- Wiley(Publisher)
Chapter Twelve ELECTROCHEMICAL CELLS AND ELECTRODE POTENTIALS “I know of nothing sublime which is not some modification of power.” —Edmund Burke Learning Objectives WHAT ARE SOME OF THE KEY THINGS WE WILL LEARN FROM THIS CHAPTER? ● Voltaic cells, p. 385 ● Using standard potentials to predict reactions, p. 387 ● Anodes, cathodes, and cell voltages (key equation: 12.19), p. 389 ● The Nernst equation (key equation: 12.22), p. 390 ● Calculating electrode potentials before and after reaction, p. 391 ● The formal potential, p. 394 An important class of titrations is reduction – oxidation or “redox” titrations, in which an oxidizing agent and a reducing agent react (see Equation 12.1). We define an oxidation as a loss of electrons to an oxidizing agent (which itself gets reduced) Oxidation is a loss of electrons. Reduction is a gain of electrons. “OIL RIG” is a good mnemonic to help remember this. 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 electrochemical cells, standard electrode potentials, the Nernst equation (which describes electrode potentials), and limitations of those potentials. Chapter 13 discusses potentiometry, the use of potential measurements for determining concentration, including the glass pH electrode and ion-selective electrodes. In Chapter 14, 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.
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