Electrochemical Systems E-Book

John Newman

• Herausgeber: John Wiley & Sons
• Kategorie: Wissenschaft und neue Technologien
• Serie: Electrochemical Society Series
• Sprache: Englisch

Seitenzahl: 1109

Veröffentlichungsjahr: 2021

Beschreibung

Provides a comprehensive understanding of a wide range of systems and topics in electrochemistry This book offers complete coverage of electrochemical theories as they pertain to the understanding of electrochemical systems. It describes the foundations of thermodynamics, chemical kinetics, and transport phenomena--including the electrical potential and charged species. It also shows how to apply electrochemical principles to systems analysis and mathematical modeling. Using these tools, the reader will be able to model mathematically any system of interest and realize quantitative descriptions of the processes involved. This brand new edition of Electrochemical Systems updates all chapters while adding content on lithium battery electrolyte characterization and polymer electrolytes. It also includes a new chapter on impedance spectroscopy. Presented in 4 sections, the book covers: Thermodynamics of Electrochemical Cells, Electrode Kinetics and Other Interfacial Phenomena, Transport Processes in Electrolytic Solutions, and Current Distribution and Mass Transfer in Electrochemical Systems. It also features three appendixes containing information on: Partial Molar Volumes, Vectors and Tensors, and Numerical Solution of Coupled, Ordinary Differential Equations. * Details fundamental knowledge with a thorough methodology * Thoroughly updated throughout with new material on topics including lithium battery electrolyte characterization, impedance analysis, and polymer electrolytes * Includes a discussion of equilibration of a charged polymer material and an electrolytic solution (the Donnan equilibrium) * A peerless classic on electrochemical engineering Electrochemical Systems, Fourth Edition is an excellent resource for students, scientists, and researchers involved in electrochemical engineering.

Leseprobe

Table of Contents

Cover

Series Title Page

Title Page

Copyright

Preface to the Fourth Edition

Preface to the Third Edition

Preface to the Second Edition

Preface to the First Edition

Reference

Chapter 1: Introduction

1.1 Definitions

1.2 Thermodynamics and Potential

1.3 Kinetics and Rates of Reaction

1.4 Transport

1.5 Concentration Overpotential and the Diffusion Potential

1.6 Overall Cell Potential

Problems

Notation

Part A: Thermodynamics of Electrochemical Cells

Chapter 2: Thermodynamics in Terms of Electrochemical Potentials

2.1 Phase Equilibrium

2.2 Chemical Potential and Electrochemical Potential

2.3 Definition of Some Thermodynamic Functions

2.4 Cell with Solution of Uniform Concentration

2.5 Transport Processes in Junction Regions

2.6 Cell with a Single Electrolyte of Varying Concentration

2.7 Cell with Two Electrolytes, One of Nearly Uniform Concentration

2.8 Cell with Two Electrolytes, Both of Varying Concentration

2.9 Lithium–Lithium Cell With Two Polymer Electrolytes

2.10 Standard Cell Potential and Activity Coefficients

2.11 Pressure Dependence of Activity Coefficients

2.12 Temperature Dependence of Cell Potentials

Problems

Notation

References

Chapter 3: The Electric Potential

3.1 The Electrostatic Potential

3.2 Intermolecular Forces

3.3 Outer and Inner Potentials

3.4 Potentials of Reference Electrodes

3.5 The Electric Potential in Thermodynamics

Notation

References

Chapter 4: Activity Coefficients

4.1 Ionic Distributions in Dilute Solutions

4.2 Electrical Contribution to the Free Energy

4.3 Shortcomings of the Debye–Hückel Model

4.4 Binary Solutions

4.5 Multicomponent Solutions

4.6 Measurement of Activity Coefficients

4.7 Weak Electrolytes

Problems

Notation

References

Chapter 5: Reference Electrodes

5.1 Criteria for Reference Electrodes

5.2 Experimental Factors Affecting Selection of Reference Electrodes

5.3 The Hydrogen Electrode

5.4 The Calomel Electrode and Other Mercury–Mercurous Salt Electrodes

5.5 The Mercury–Mercuric Oxide Electrode

5.6 Silver–Silver Halide Electrodes

5.7 Potentials Relative to a Given Reference Electrode

Notation

References

Chapter 6: Potentials of Cells with Junctions

6.1 Nernst Equation

6.2 Types of Liquid Junctions

6.3 Formulas for Liquid‐Junction Potentials

6.4 Determination of Concentration Profiles

6.5 Numerical Results

6.6 Cells with Liquid Junction

6.7 Error in the Nernst Equation

6.8 Potentials Across Membranes

6.9 Charged Membranes Immersed in an Electrolytic Solution

Problems

Notation

References

Part B: Electrode Kinetics and other Interfacial Phenomena

Chapter 7: Structure of the Electric Double Layer

7.1 Qualitative Description of Double Layers

7.2 Gibbs Adsorption Isotherm

7.3 The Lippmann Equation

7.4 The Diffuse Part of the Double Layer

7.5 Capacity of the Double Layer in the Absence of Specific Adsorption

7.6 Specific Adsorption at an Electrode–Solution Interface

Problems

Notation

References

Chapter 8: Electrode Kinetics

8.1 Heterogeneous Electrode Reactions

8.2 Dependence of Current Density on Surface Overpotential

8.3 Models for Electrode Kinetics

8.4 Effect of Double‐Layer Structure

8.5 The Oxygen Electrode

8.6 Methods of Measurement

8.7 Simultaneous Reactions

Problems

Notation

References

Chapter 9: Electrokinetic Phenomena

9.1 Discontinuous Velocity at an Interface

9.2 Electro‐Osmosis and the Streaming Potential

9.3 Electrophoresis

9.4 Sedimentation Potential

Problems

Notation

References

Chapter 10: Electrocapillary Phenomena

10.1 Dynamics of Interfaces

10.2 Electrocapillary Motion of Mercury Drops

10.3 Sedimentation Potentials for Falling Mercury Drops

Notation

References

Part C: Transport Processes in Electrolytic Solutions

Chapter 11: Infinitely Dilute Solutions

11.1 Transport Laws

11.2 Conductivity, Diffusion Potentials, and Transference Numbers

11.3 Conservation of Charge

11.4 The Binary Electrolyte

11.5 Supporting Electrolyte

11.6 Multicomponent Diffusion by Elimination of the Electric Field

11.7 Mobilities and Diffusion Coefficients

11.8 Electroneutrality and Laplace'S Equation

11.9 Moderately Dilute Solutions

Problems

Notation

References

Chapter 12: Concentrated Solutions

12.1 Transport Laws

12.2 The Binary Electrolyte

12.3 Reference Velocities

12.4 The Potential

12.5 Connection with Dilute‐Solution Theory

12.6 Example Calculation Using Concentrated Solution Theory

12.7 Multicomponent Transport

12.8 Liquid‐Junction Potentials

Problems

Notation

References

Chapter 13: Thermal Effects

13.1 Thermal Diffusion

13.2 Heat Generation, Conservation, and Transfer

13.3 Heat Generation at an Interface

13.4 Thermogalvanic Cells

13.5 Concluding Statements

Problems

Notation

References

Chapter 14: Transport Properties

14.1 Infinitely Dilute Solutions

14.2 Solutions of a Single Salt

14.3 Mixtures of Polymers and Salts

14.4 Types of Transport Properties and Their Number

14.5 Integral Diffusion Coefficients for Mass Transfer

Problem

Notation

References

Chapter 15: Fluid Mechanics

15.1 Mass and Momentum Balances

15.2 Stress in a Newtonian Fluid

15.3 Boundary Conditions

15.4 Fluid Flow to a Rotating Disk

15.5 Magnitude of Electrical Forces

15.6 Turbulent Flow

15.7 Mass Transfer in Turbulent Flow

15.8 Dissipation Theorem for Turbulent Pipe Flow

Problem

Notation

References

Part D: Current Distribution and Mass Transfer in Electrochemical Systems

Chapter 16: Fundamental Equations

16.1 Transport in Dilute Solutions

16.2 Electrode Kinetics

Notation

Chapter 17: Convective‐Transport Problems

17.1 Simplifications for Convective Transport

17.2 The Rotating Disk

17.3 The Graetz Problem

17.4 The Annulus

17.5 Two‐Dimensional Diffusion Layers in Laminar Forced Convection

17.6 Axisymmetric Diffusion Layers in Laminar Forced Convection

17.7 A Flat Plate in a Free Stream

17.8 Rotating Cylinders

17.9 Growing Mercury Drops

17.10 Free Convection

17.11 Combined Free and Forced Convection

17.12 Limitations of Surface Reactions

17.13 Binary and Concentrated Solutions

Problems

Notation

References

Chapter 18: Applications of Potential Theory

18.1 Simplifications For Potential‐Theory Problems

18.2 Primary Current Distribution

18.3 Secondary Current Distribution

18.4 Numerical Solution by Finite Differences

18.5 Principles of Cathodic Protection

Problems

Notation

References

Chapter 19: Effect of Migration on Limiting Currents

19.1 Analysis

19.2 Correction Factor for Limiting Currents

19.3 Concentration Variation of Supporting Electrolyte

19.4 Role of Bisulfate Ions

19.5 Paradoxes with Supporting Electrolyte

19.6 Limiting Currents for Free Convection

Problems

Notation

References

Chapter 20: Concentration Overpotential

20.1 Definition

20.2 Binary Electrolyte

20.3 Supporting Electrolyte

20.4 Calculated Values

Problems

Notation

References

Chapter 21: Currents Below the Limiting Current

21.1 The Bulk Medium

21.2 The Diffusion Layers

21.3 Boundary Conditions and Method of Solution

21.4 Results for the Rotating Disk

Problems

Notation

References

Chapter 22: Porous Electrodes

22.1 Macroscopic Description of Porous Electrodes

22.2 Nonuniform Reaction Rates

22.3 Mass Transfer

22.4 Battery Simulation

22.5 Double‐Layer Charging and Adsorption

22.6 Flow‐Through Electrochemical Reactors

Problems

Notation

References

Chapter 23: Semiconductor Electrodes

23.1 Nature of Semiconductors

23.2 Electric Capacitance at the Semiconductor–Solution Interface

23.3 Liquid‐Junction Solar Cell

23.4 Generalized Interfacial Kinetics

23.5 Additional Aspects

Problems

Notation

References

Chapter 24: Impedance

24.1 Frequency Dispersion at a Disk Electrode

24.2 Modulated Flow With a Disk Electrode

24.3 Porous Electrodes for Batteries

24.4 Kramers–Kronig Relation

Problems

Notation

References

Appendix A: Partial Molar Volumes

Appendix B: Vectors and Tensors

Reference

Appendix C: Numerical Solution of Coupled, Ordinary Differential Equations

C.1 Errors in Finite‐Difference Calculations

C.2 Convergence Over Nonlinearities

C.3 Solution of Coupled, Linear, Difference Equations

C.4 Program for Coupled, Linear Difference Equations

C.5 Program for the Effect of Ionic Migration on Limiting Currents

C.6 Second Example: Multicomponent Diffusion

C.7 Discussion and Conclusions

References

Index

End User License Agreement

List of Tables

Chapter 2

TABLE 2.1 Effect of solubility of silver chloride for decreasing values of bulk HCl concentration

TABLE 2.2 Selected standard electrode potentials referred to the hydrogen electrode in aqueous solutions at 25°C

TABLE 2.3 Additional standard electrode potentials in aqueous solutions at 25°C

TABLE 2.4 Thermodynamic data for the hydrogen/oxygen fuel cell evaluated at standard reference conditions of 298.15 K and 1 bar for liquid water and the ideal‐gas state for gaseous species

Chapter 4

TABLE 4.1 Debye–Hückel parameters for aqueous solutions

TABLE 4.2 Values of

β

(kg/mol) for 1–1 electrolytes at 25°C and for

Ba

= 1 (kg/mol)

1/2

TABLE 4.3 Values of

β

and

Ba

for 2–1 and 1–2 electrolytes at 25°C

TABLE 4.4 Values of

β

for ions of like charge

Chapter 6

TABLE 6.1 Values of ΔΦ for various junctions and various models at 25°Ca

TABLE 6.2 Values of ΔΦ for a Ag–AgCl electrode in HCl solutions at 25°Ca

Chapter 7

TABLE 7.1 Potential of zero charge for mercury (relative to a normal calomel electrode in KCl) for various electrolytic solutions at 25°C

Chapter 8

TABLE 8.1 Free energy of surface species at the open‐circuit potential of the oxygen electrode (

bar,

mol/L), 1.229 V, with respect to the standard H

2

electrode

Chapter 9

TABLE 9.1 Dimensionless flow rate

μ

〈

〉/

λq

2

E

z

as function of

R

0

in the absence of a pressure drop

Chapter 11

TABLE 11.1 Values of equivalent conductances and diffusion coefficients of selected ions at infinite dilution in water at 25°C

Chapter 12

TABLE 12.1 Comparison of results for binary electrolytes

Chapter 14

TABLE 14.1 Transport properties and their numbera

Chapter 17

TABLE 17.1 Eigenvalues and coefficients for the Graetz series

TABLE 17.2 Coefficient

C

expressing the rate of mass transfer for free convection at a vertical plate from a binary fluid with a uniform density difference between the vertical surface and the bulk solution

Chapter 18

TABLE 18.1 Supplemental potential map for the base case

TABLE 18.2 Design spreadsheet

Chapter 20

TABLE 20.1 Values of concentration overpotential

η

c

(in mV) for copper deposition on a rotating disk from solutions of copper sulfate and sulfuric acid, with complete dissociation of bisulfate ions

TABLE 20.2 Values of concentration overpotential

η

c

(in mV) for reduction of ferricyanide ions on a rotating disk from solutions equimolar in potassium ferricyanide and potassium ferrocyanide and with various amounts of added potassium hydroxide

Chapter 22

TABLE 22.1 Operating conditions, design results, and costs for removal of lead and copper ions from given solutions

Appendix B

TABLE B.1 Vector and tensor algebra and calculus

List of Illustrations

Chapter 1

Figure 1.1 Volta's first battery comprised of a sandwich of zinc with its oxide layer, salt solution, and silver with its oxide layer. While the original Volta pile used an electrolyte of NaCl in water, modern batteries use aqueous KOH to increase the conductivity and the concentration of OH

−

.

Figure 1.2 Schematic of the relative energy of the electron in reduction and oxidation reactions. During a reduction reaction, electrons are transferred from the electrode to the lowest unoccupied energy level of a reactant species. During oxidation, electrons are transferred from the highest occupied energy level of the reactant to the electrode.

Figure 1.3 Dependence of current density on surface overpotential at 25°C.

Figure 1.4 Tafel plot of the relationship between current density and surface overpotential at 25°C.

Figure 1.5 Two concentric copper electrodes with the annulus filled with electrolyte. The inner electrode can be rotated.

Figure 1.6 Distribution of the potential in solution between cylindrical electrodes.

Figure 1.7 Concentration profile in the annular space between the electrodes. The dashed curve refers to the absence of a radial component of velocity. The solid curve refers to the presence of turbulent mixing.

Figure 1.8 Streamlines for free convection in the annular space between two cylindrical electrodes.

Figure 1.9 Concentration cell.

Figure 1.10 Placement of reference electrodes (1, 2, and 3) in the solution between cylindrical electrodes. The concentration profile shown corresponds to turbulent mixing at a current somewhat below the limiting current.

Figure 1.11 Concentration overpotentials at a cathode in 0.1

M

CuSO

4

.

Figure 1.12 The dependence of the cell potential and its component overpotentials on current for concentric cylinders, the inner of which rotates. The overpotentials for the anode are small for this particular system and are not shown.

Figure 1.13 Current–potential relations with sulfuric acid added as a supporting electrolyte.

Chapter 2

Figure 2.1 Variation of the molar activity coefficient of aqueous acetic acid with concentration. For

1 + 

d

 ln 

f

+−

/

d

 ln 

c

, divide the ordinate scale by 2.

Figure 2.2 (a) Schematic of a concentration cell with two lithium electrodes and a junction between two polymer electrolytes with different salt concentrations (molalities) in contact with each other. (b) Open‐circuit potential

U

, as a function of molality

m

, of PEO/LiTFSI, with a reference molality of 1.36 mol/kg, measured before diffusion substantially changes the electrolyte concentration at the electrodes.

Figure 2.3 Simplified Pourbaix diagram of potential

versus

pH for zinc at 1 mol/kg concentration of Zn

2+

and

, showing regions of stability of ZnO and

. For reference, the dashed lines for the evolution of oxygen and hydrogen indicate the limits of stability of water.

Figure 2.4 Open‐circuit potential and enthalpy potential for a hydrogen–oxygen fuel cell, using either liquid or gaseous water as the product, with all pressures at 1 bar.

Chapter 3

Figure 3.1 Normal components of the electric field at an interface. The interface may have a charge

σ

per unit area.

Figure 3.2 Potential difference between two metal spheres for an average surface charge of 10 μC/cm

2

.

Figure 3.3 Intermolecular potential energy for two ions at a distance

r

apart.

Figure 3.4 Intermolecular force between two ions.

Figure 3.5 Movement of a charged particle from a cavity in one phase to a cavity in another phase. This thought experiment is used to define the Volta potential and the contact potential difference between two metals.

Figure 3.6 Use of reference electrodes to investigate the potential in a solution. Silver–silver chloride reference electrodes are represented by

α

and

β

. The vessel on the left also contains two working electrodes.

Figure 3.7 Use of reference electrodes to investigate potential variations within a solution.

Chapter 4

Figure 4.1 Ionic distributions near a central cation, according to the theory of Debye and Hückel, for a 0.1

M

aqueous solution of a uni‐univalent electrolyte at 25°C.

Figure 4.2 Mean molal activity coefficients of HCl (from Ref. [13]) and HNO

3

(from Ref. [23]) and the ratio of the activity coefficients of the two acids.

Figure 4.3 Ratio of the activity coefficients of HCl and HNO

3

plotted against the molality.

Figure 4.4 Correlation of the second dissociation constant of sulfuric acid with the true ionic strength.

Chapter 6

Figure 6.1 Calculated values of ΔΦ for free‐diffusion, restricted‐diffusion, and continuous‐mixture boundaries between HCl and KCl. (a–c) graphs are for given concentrations of KCl on one side of the boundary. (d–f) graphs are for a given ratio of concentrations on the two sides of the boundary. The dashed lines represent ideal‐solution calculations; the solid lines include activity‐coefficient corrections.

Figure 6.2 Schematic of a charged crosslinked polymer membrane with negative charges covalently bound to the polymer chains in contact with an electrolytic solution. Some of the ions in the electrolytic solution enter the membrane.

Figure 6.3 Dependence of the coion and counterion concentrations in the membrane on concentration of the electrolytic solution for different values of

v

−

/

v

+

. Ion concentrations in the membrane are normalized by the charge concentration within the membrane, assumed to be negatively charged.

f

m

has been taken to be unity.

Figure 6.4 Four electrodes in a cell with a liquid junction.

Chapter 7

Figure 7.1 Solid–solution interface with no charge in the solid.

Figure 7.2 Excess electric charge density in the diffuse part of the double layer.

Figure 7.3 Dipole moment in a water molecule.

Figure 7.4 Oriented water molecules at an interface with no charge in the solid.

Figure 7.5 Metal–solution interfaces arranged so that the charge on the metal can be varied. Now there is a charge in the metal near the interface with the solution.

Figure 7.6 Steady potential distribution in a system of ideally polarizable electrodes.

Figure 7.7 Structure of the double layer. The charge on the metal side of the interface is

q

. Specifically adsorbed ions or molecules are located at the inner Helmholtz plane, while solvated adsorbed ions are located beyond (but not quite at) the outer Helmholtz plane. The diffuse layer is like the bulk of the solution except that it is not electrically neutral, but rather has a net charge

q

2

. The diffusion layer is electrically neutral but may have a nonuniform salt concentration.

Figure 7.8 Apparatus for determining the point of zero charge on mercury in an electrolytic solution.

Figure 7.9 Apparatus for charging mercury drops in an electrolytic solution.

Figure 7.10 Interfacial, nonhomogeneous region of thickness

τ

between two homogeneous phases.

Figure 7.11 System for applying a potential to an ideally polarizable electrode.

Figure 7.12 Interfacial tension of mercury as a function of potential for several electrolytic solutions at 18°C. Potentials relative to a normal calomel electrode are shifted by 0.48 V. These are referred to as electrocapillary curves because the surface tension is often measured with a capillary electrometer.

Figure 7.13 Charge and adsorption of sodium and chloride ions at a mercury interface in contact with 0.3 

M

NaCl at 25°C. The surface concentrations of the ions are expressed as

z

i

F

Γ

i

.

Figure 7.14 Charge and adsorption of sodium and chloride ions at a mercury interface in contact with a 1 

M

NaCl solution at 25°C. The surface concentrations of the ions are expressed as

z

i

F

Γ

i

.

Figure 7.15 Double‐layer capacity for mercury in contact with NaCl solutions at 25°C. Potentials are relative to the electrocapillary maximum.

Figure 7.16 Double‐layer capacity for mercury in contact with NaF solutions at 25°C. Potentials are relative to the normal calomel electrode.

Chapter 8

Figure 8.1 Current–potential relation for an electrode exhibiting passivation.

Figure 8.2 A 3‐valent ion (ferricyanide) from the outer Helmholtz plane can be adsorbed at the inner Helmholtz plane, where it can react with an electron from the metal electrode. The resulting 4‐valent ion (ferrocyanide) is then desorbed.

Figure 8.3 Active‐intermediate diagram, showing the potential energy of an electrode–ion system as a function of distance of the electron from the ion.

Figure 8.4 Potential–energy diagram for an elementary charge‐transfer step. The solid curve is for

V

=

V

1

. The dashed curve is for

V

=

V

2

, where

V

2

is greater than

V

1

.

Figure 8.5 Anodic and cathodic contributions to the current density (from equation 8.35 with

β

3

= 0.5) plotted against the potential relative to a copper electrode in 1

M

CuSO

4

(see Section 5.7).

Figure 8.6 Tafel plot of surface overpotential for the hydrogen‐evolution reaction on various electrode materials.

Figure 8.7 Tafel plot of overpotentials for the hydrogen electrode when the Heyrovský reaction (H

+

 + e

−

 + H

ad

 → H

2

) is the rate‐determining step and the Volmer reaction is rapid.

Figure 8.8 Map of the relative rates for the three elementary steps involved in the hydrogen electrode reaction. The three intersecting lines give the directions for increasing rate constants for the Volmer (V), Heyrovský (H), and Tafel (T) reactions. The other arrows indicate where the six limiting cases can be found, the first letter standing for the rate‐determining step and the second letter showing the equilibrated step.

Figure 8.9 The potential‐dependent surface free energies of surface species in the oxygen electrode on RuO

2

.

Figure 8.10 Calculated current–potential curves for oxygen reduction and evolution on a Tafel plot. The values for RuO

2

use equilibrium constants slightly modified from those in Ref. [22] (to make the open‐circuit potential agree with that in Table 2.2). The ideal system uses unity values for

K

1

,

K

2

,

K

3

, and

K

4

, thereby giving a fractional coverage of 0.25 for each of the four surface species. This figure is very similar to Figure 8.7, but with the axes rotated by 90°.

Figure 8.11 Theoretical polarographic curves for the reduction of oxygen, with neglect of the anodic reaction terms. In this,

K

=

D

B

k

c

2

/

D

A

k

c

3

.

Chapter 9

Figure 9.1 Velocity produced by a tangential electric field in the diffuse charge layer. A positive charge in the diffuse layer will produce a negative zeta potential and will result in a positive value of

if

E

x

is also positive.

Figure 9.2 Velocity profile in the capillary when there is no pressure drop; this is the case of electro‐osmosis.

Figure 9.3 Velocity profile in the capillary when there is no net fluid flow.

Chapter 11

Figure 11.1 Accumulation due to differences in the fluxes at the faces of a volume element.

Chapter 12

Figure 12.1 Potentials of Pb, PbO

2

, and Hg/Hg

2

SO

4

electrodes versus the molality of sulfuric acid, plotted in such a way that vertical distances give the potenial differences between electrodes even if the composition is different.

Chapter 13

Figure 13.1 Thermogalvanic cell.

Figure 13.2 Thermocouple.

Chapter 14

Figure 14.1 Multicomponent diffusion coefficients of KC1–H

2

O at 25°C.

Figure 14.2 Empirical function

G

for various systems.

Figure 14.3 Diffusion coefficient of chloride ion in various aqueous solutions at 25°C.

Figure 14.4 Diffusion coefficient of chloride ion with a viscosity factor.

Figure 14.5 Lithium ion diffusion coefficient in lithium chloride solutions at various temperatures.

Figure 14.6 Chloride ion diffusion coefficient in lithium chloride solutions at various temperatures.

Figure 14.7 Chemical formulae and schematic depictions of (a) homopolymer and (b) block copolymer electrolytes. The arrows in (a) and (b) represent one of the possible pathways for salt diffusion in the homopolymer and block copolymer electrolytes, respectively. In the block copolymer, diffusion is limited to the bright phase. (c) An electron micrograph of a block copolymer electrolyte: polystyrene‐

b

‐poly(ethylene oxide) with a lithium salt. The conducting domains appear bright in the micrograph.

Figure 14.8 Complete characterization of ion transport in a polymer electrolyte (PEO/LiTFSI) at 90°C. (a) Conductivity,

κ

; (b) salt diffusion coefficient,

D

; (c) cation transference number,

; and (d) the thermodynamic factor as a function of molality,

m

. The approximate transference number based on the assumption of an ideal electrolyte,

t

+,id

, is also shown in (c).

Figure 14.9 (a–c) Stefan–Maxwell diffusion coefficients of PEO/LiTFSI electrolytes as a function of molality. (d) Reciprocal of the Stefan–Maxwell diffusion coefficients in (b) and (c) as a function of molality.

Figure 14.10 Concentration profiles in PEO/LiTFSI electrolytes predicted using characterization data shown in Figure 14.8. Curves depend on the product of steady‐state current density (

i

ss

) and electrode separation (

L

), given at the top of parts a–c. Same scale applies to all figures.

Figure 14.11 Potential profiles in PEO/LiTFSI electrolytes predicted using characterization data shown in Figure 14.8. Curves depend on the product of steady‐state current density (

i

ss

) and electrode separation (

L

), given at the top of parts a–c. Same scale applies to all figures. The inset in (a) shows curves on an expanded scale for clarity.

Figure 14.12 Characteristics of lithium‐polymer‐lithium cells with

L

= 500 μm (thickness of the polymer electrolyte) for a constant steady‐state current,

i

ss

= 0.02 mA/cm

2

. The cell potential,

φ

ss

, normalized by

L

, is plotted as a function of average molality,

m

av

. The circles represent experimental measurements. The solid curve represents theoretical predictions based on the characterization data in Figure 14.8 and equations 12.40 and 12.44.

Figure 14.13 (a) Schematic of a composite electrolyte with a lamellar morphology with many randomly oriented grains sandwiched between two electrodes. (b) Schematic of a grain showing the salt localized in one of the lamellae.

Chapter 15

Figure 15.1 Tangential forces on an interfacial element lying in the

x

,

z

plane.

Figure 15.2 Velocity profiles for a rotating disk.

Figure 15.3 Universal velocity profile for fully developed turbulent flow.

Figure 15.4 Representation of the eddy viscosity as a “universal” function of the distance from the wall.

Figure 15.5 Variation of the eddy diffusivity near a wall for fully developed turbulent flow.

Figure 15.6 The eddy‐viscosity profiles of Nikuradse for his lowest 4 Reynolds numbers. The dashed line is the limit curve for large Reynolds numbers. Points for the lower Reynolds numbers generally lie slightly higher than the limit curve. Here, the Reynolds numbers are 4000, 6100, 9200, and 16,700.

Chapter 17

Figure 17.1 Concentration profile in the diffusion layer.

Figure 17.2 Rotating‐disk electrode.

Figure 17.3 Dimensionless mass‐transfer rates for a rotating disk.

Figure 17.4 Graetz functions.

Figure 17.5 Dimensionless cup‐mixing concentration difference Θ

m

and the local Nusselt number (divided by Lévêque's solution). For comparison with the latter, the corresponding form of the Lévêque series is shown for two and three terms.

Figure 17.6 Coefficient for mass transfer in annuli.

Figure 17.7 Current distribution on planar electrodes.

Figure 17.8 Plane electrodes in the walls of a flow channel.

Figure 17.9 Electrode on an axisymmetric body with axisymmetric flow.

Figure 17.10 Sketch of Taylor vortices.

Figure 17.11 Photograph of Taylor vortices at a Reynolds number of 143 with

r

0

/

r

i

= 1.144

.

Chapter 18

Figure 18.1 Two plane electrodes opposite each other in the walls of an insulating flow channel, showing equipotential surfaces (‐ ‐ ‐) and current lines (—).

Figure 18.2 Current distribution on planar electrodes. Here

x

is measured from the edge of the electrode, not the center.

Figure 18.3 Behavior of the primary current distribution near the edge of an electrode.

Figure 18.4 Current (‐ ‐ ‐) and potential (—) lines for a disk electrode.

Figure 18.5 Secondary current distribution for linear polarization at a disk electrode.

Figure 18.6 Secondary current distribution for Tafel polarization at a disk electrode.

Figure 18.7 Current density at the center of the disk when concentration polarization is absent.

Figure 18.8 Primary current distribution and potential distribution for a uniform current density on a disk electrode.

Figure 18.9 Simplified Pourbaix diagram for iron in water at 25°C. The oxide phases Fe

2

O

3

and Fe

3

O

4

are not shown.

Figure 18.10 Cathodic protection system.

Figure 18.11 Potential relative to an adjacent saturated Cu/CuSO

4

reference electrode. Arrows indicate direction of pH shift induced by the expected electrochemical reactions.

Figure 18.12 Protected pipeline and parallel, cylindrical anode that provides the required current.

Figure 18.13 Current–potential curve for a local element. The electrode potential is the potential of the protected surface minus that of the adjacent soil, as assessed with a Cu/CuSO

4

reference electrode. The current density for oxygen is shown constant at −1, corresponding to the limiting current density in this range of potentials.

Figure 18.14 Equipotential contours (of the quantity

κ

Φ/

i

avg

r

c

)

for the cathodic protection system with

r

c

/

r

a

= 24

and

d

/

r

c

= 8

and for a uniform current density on the cathode.

Figure 18.15 Potentials in the conductors of the system. Variable resistors are suggested between the anode and the power supply cable.

Figure 18.16 Variation of potential in the soil around the cathode and the average potential drop in the soil between the anode and the cathode. Note the logarithmic scales. For comparison, the points show the potential difference between the anode and the cathode when the soil potential near the cathode is uniform (corresponding to the primary current distribution in this system).

Figure 18.17 Equipotential contours for two symmetrically placed anodes, with

r

c

/

r

a

= 24

and

d

/

r

c

= 1

.

Figure 18.18 Equipotential contours for one anode,

r

c

/

r

a

= 24

and

d

/

r

c

= 1

.

Figure 18.19 Variation of potential in the soil around the cathode for one or two anodes. Asymptotes of values of

κΔ

Φ/

i

avg

r

c

for large

d

/

r

c

are also shown (dashed lines).

Figure 18.20 Correction for the potential difference in the soil between the anode and the near part of the cathode, for a cylindrical cathode with a single anode.

Figure 18.21 Correction for the potential difference in the soil between the anode and the near part of the cathode, for a cylindrical cathode with two anodes symmetrically placed.

Figure 18.22 Ring–disk electrode, frequently rotated to provide a known hydrodynamic flow.

Figure 18.23 A section of a plate of thickness

L

with a two‐dimensional slot of width

h

through its thickness.

Figure 18.24 An electrode at angle

θ

= 0 meets an insulator at angle

θ = α

.

Figure 18.25 Top view of the Hull cell.

Chapter 19

Figure 19.1 Effect of migration on limiting currents for metal deposition on a disk electrode.

Figure 19.2 Effect of migration on limiting currents in discharge of hydrogen ions from KCl solutions. Lines represent values calculated with the present theory.

Figure 19.3 Effect of migration on limiting currents for a redox reaction. Equimolar potassium ferrocyanide and ferricyanide in KOH, for a disk electrode.

Figure 19.4 Concentration difference of the added ion divided by that of the reactant. The abscissa scale is defined in Figures 19.1 and 19.2.

Figure 19.5 Surface concentrations for the anodic reaction in the K

3

Fe(CN)

6

–K

4

Fe(CN)

6

–KOH system.

Figure 19.6 Surface concentrations for the cathodic reaction in the K

3

Fe(CN)

6

–K4Fe(CN)

6

–KOH system.

Figure 19.7 Conductivity of aqueous solutions of copper sulfate and sulfuric acid at 25°C.

Figure 19.8 Effect of migration in the CuSO

4

–H

2

SO

4

system with no dissociation and with complete dissociation of bisulfate ions.

Figure 19.9 Effect of migration for a rotating‐disk electrode.

Figure 19.10 Effect of migration for a growing mercury drop or in a stagnant diffusion cell.

Figure 19.11 Surface concentration change for a rotating‐disk electrode.

Figure 19.12 Surface concentration change for a growing mercury drop or in a stagnant diffusion cell.

Figure 19.13 Effect of migration in a Nernst diffusion layer.

Figure 19.14 Surface concentration change in a Nernst diffusion layer.

Figure 19.15 Concentration differences of sulfuric acid possible in the copper sulfate, sulfuric acid system with complete and with no dissociation of bisulfate ions and for several hydrodynamic situations.

Figure 19.16 Coefficients for shear stress (complete dissociation only) and mass transfer in the CuSO

4

–H

2

SO

4

system. Dashed curves show for comparison values of

I

L

/

I

D

for the rotating disk.

Figure 19.17 Velocity profiles for binary salt solution (CuSO

4

) and for CuSO

4

with excess H

2

SO

4

(

r

= 0.99998) completely dissociated and undissociated.

Figure 19.18 Coefficient for mass‐transfer rate in the supported ferricyanide–ferrocyanide systems, for equal bulk concentrations of K

3

Fe(CN)

6

and K

4

Fe(CN)

6

.

Figure 19.19 Surface concentrations in the supported ferricyanide–ferrocyanide systems, for equal bulk concentrations of K

3

Fe(CN)

6

and K

4

Fe(CN)

6

.

Figure 19.20 Velocity profiles for various values of

r

for cathodic reduction of ferricyanide ions with KOH supporting electrolyte.

Figure 19.21 Normalized density profiles for binary salt solution (CuSO

4

), for CuSO

4

with excess H

2

SO

4

(

r

= 0.99998), and for equimolar ferricyanide–ferrocyanide with excess KOH (cathodic reaction,

).

Chapter 21

Figure 21.1 Surface concentration for Tafel kinetics.

Figure 21.2 Current distribution for Tafel kinetics with an appreciable fraction of the limiting current.

Figure 21.3 Current density at the center of the disk.

Figure 21.4 Overpotentials for copper deposition on a rotating disk. Dashed line is ohmic drop for the primary current distribution;

, and

η

s

are evaluated at the center of the disk.

Figure 21.5 Total overpotential at several positions on the disk.

Figure 21.6 Channel flow cell, with diffusion layers shown for the two electrodes.

Chapter 22

Figure 22.1 Schematic of a one‐dimensional porous electrode.

Figure 22.2 Electric analog of a porous electrode with ohmic resistances representing matrix and pore solution (upper and lower horizontal resistors) and kinetic resistance (vertical elements). The vertical branches also have elements representing a cell of potential

U

.

Figure 22.3 Reduced current distribution for Tafel polarization with

σ

= κ.

Figure 22.4 Potential distributions for Tafel polarization with

σ

= κ. Here

β

=

α

a

F

/

RT

(or −

α

c

F

/

RT

for cathodic currents).

Figure 22.5 Potential of the metal backing plate as it depends on

δ

. Here

β

= 

α

a

F

/

RT

(or −

α

c

F

/

RT

for cathodic currents).

Figure 22.6 Pictorial of the battery system.

Figure 22.7 Equivalent‐circuit representation for the reaction‐zone model.

Figure 22.8 Specific energy for various electrode thicknesses and porosities. The optimum is

ε = 0.227 and

L

+

/

L

s

= 1.95

.

Figure 22.9 Optimum electrode thickness as it depends on the parameter determined by discharge time, electrode capacity density, open‐circuit potential, and separator parameters.

Figure 22.10 Optimum porosity as it depends on the parameter determined by discharge time, electrode capacity density, open‐circuit potential, and separator parameters.

Figure 22.11 Specific energy versus average specific power. For three of the curves, the electrode thickness and porosity are optimized at the values of

T

given. For the envelope curve, the electrode thickness and porosity are optimized for each point on the curve. Here

ρ

+

/

ρ

s

=

b

= 1.

Figure 22.12 Schematic diagram of the LiAl–FeS cell, as an example of the cell‐sandwich model.

Figure 22.13 Position dependence of mole fraction of LiCl at different discharge times, for

X

‐phase mechanism. Dashed line represents saturation limit for LiCl at 450°C.

Figure 22.14 Composition profile through the cell sandwich, with

J

‐phase mechanism.

Figure 22.15 Comparison of theoretical and experimental discharge curves for the

X

‐phase mechanism.

Figure 22.16 Discharge curve with

J

‐phase mechanism.

Figure 22.17 Comparison of model and experimental results for the potential of the FeS

2

electrode in a LiAl–FeS

2

cell, relative to a LiAl (

α

 − 

β

) reference electrode (450°C, 50 mA/cm

2

,

). The reversible, thermodynamic potential is also shown in order to display more clearly the losses of the system.

Figure 22.18 Volume fraction of solid phases and electrolyte in the positive electrode of LiAl–FeS cell discharging by the

X

‐phase mechanism.

Figure 22.19 Comparison of experimental and theoretical results for potentiostatic double layer charging of porous PbO

2

electrodes. Area = 241 cm

2

,

L

= 0.095 cm, temperature = 28°C. ▪: freshly prepared PbO

2

electrode, Δ

V

= 2.51 mV,

λ

= 1.33,

aC

= 23.33 F/cm

3

; •: cycled PbO

2

electrode, Δ

V

= 1.52 mV,

λ

= 0.768,

aC

= 26 F/cm

3

. (

λ

is a ratio of the external resistance, Ω cm

2

, to

L

/

κ

.) Solid curves are theoretical.

Figure 22.20 Two configurations of a flow‐through porous electrode showing the placement of the counterelectrode (CE) and the current collector (CC). A desirable configuration is to have the CE upstream, as in (a), but the downstream CE in (b) may keep reaction products from the CE out of the working electrode. A third possibility (not shown) is to place the CE along the working electrode.

Chapter 23

Figure 23.1 Density of states

N

(

E

) versus energy level near the band edges. More, unfilled bands would lie above the energies shown, and more, filled bands would lie below.

Figure 23.2 Activity coefficients of holes and electrons.

Figure 23.3 Junction between two metals.

Figure 23.4 Junction between semiconductors of different doping levels.

Figure 23.5 Potential, field, and charge density. Dopant levels are 2 × 10

16

and 6 × 10

16

/cm

3

in the p‐ and n‐doped regions, respectively.

T

= 300 K, and

E

g

= 1.4 J/C.

Figure 23.6 Capacity of the space‐charge region.

Figure 23.7 Capacity in a Mott–Schottky plot.

Figure 23.8 Sketch of the liquid‐junction photovoltaic cell, showing assumed interfacial reactions.

Figure 23.9 Potential distribution for the photoelectrochemical cell with no interfacial kinetic limitations. Curve (a), open circuit in the dark; curve (b), open circuit under 882 W/m

2

illumination; and curve (c), near short circuit (

i

= −23.1 mA/cm

2

) under illumination. The semiconductor is

n

‐GaAs, and the solution contains 0.8

M

K

2

Se, 0.1

M

K

2

Se

2

, and 1

M

KOH.

Figure 23.10 Concentration distributions for the photoelectrochemical cell with no interfacial kinetic limitations. Curves and conditions are as in Figure 23.9. Concentrations are made dimensionless with the net dopant concentration

N

d

 − 

N

a

.

Figure 23.11 Computed current–potential curves for an

n

‐type GaAs anode with dopant concentration as a parameter. The semiconductor is 10 Debye lengths thick in each case.

Figure 23.12 Planes in the interfacial region between an electrode and a solution. Species at plane

α,

which is closer to the electrode terminal, can react with species at plane

δ

, which is closer to the solution. Reaction can include transfer of an ion or molecule from plane

δ

to plane

α

.

Figure 23.13 Variation of cavity potential Φ from a platinum counterelectrode, through an electrolytic solution and a semiconductor electrode, and into a platinum current collector. Potential jumps occur at the phase boundaries, but the overall cell potential is zero at open circuit in the dark.

Chapter 24

Figure 24.1 Equivalent circuit with, from left to right, a power source, a rectifier, a capacitor, a resistor, an inductor, another capacitor, and a DC output.

Figure 24.2 Wave forms, with, from top to bottom, a 60‐Hz sinusoidal alternating current (AC), direct current (DC) after half‐wave rectification, DC after full‐wave rectification, and DC at the outlet after some filtering.

Figure 24.3 Equivalent circuit of the interface, showing a resistor for electrochemical kinetics in parallel with a capacitor for the double‐layer capacity.

Figure 24.4 Frequency dependence of apparent capacity on a smooth disk in the absence of faradaic reactions.

Figure 24.5 Comparison between the theoretical curve from equation 24.23 and the experimental data in galvanostatic mode. The data were obtained with the

system at

. Supporting electrolyte is 1 

M

KCl, Sc = 1200,

. (a) phase shift versus dimensionless frequency, (b) normalized amplitude versus dimensionless frequency.

Figure 24.6 Dependence of the impedance response on the diffusion coefficient

D

s

in the positive‐electrode material Li

y

TiS

2

.

Appendix C

Figure C.1 BANDmap for migration program set up for oxygen reduction from a solution of NaCl. An alternative would be to use bulk boundary conditions only for the first four equations (for the potential and three concentrations) and to let one ion concentration in the bulk, say Cl

−

, to be determined by electroneutrality.

Figure C.2 BANDmap for the effect of migration, using multicomponent diffusion equations. Reduction of O

2

from a solution of NaCl.

Figure C.3 Concentration profiles for reduction of O

2

from a solution of NaCl, as calculated by the transient multicomponent‐diffusion program.

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THE ELECTROCHEMICAL SOCIETY SERIES

The Electron Microprobe

Edited by T. D. McKinley, K. F. J. Heinrich, and D. B. Wittry

Chemical Physics of Ionic Solutions

Edited by B. E. Conway and R. G. Barradas

High‐Temperature Materials and Technology

Edited by Ivor E. Campbell and Edwin M. Sherwood

Alkaline Storage Batteries

S. Uno Falk and Alvin J. Salkind

The Primary Battery (in Two Volumes)

Volume I

Edited by George W. Heise and N. Corey Cahoon

Volume II

Edited by N. Corey Cahoon and George W. Heise

Zinc‐Silver Oxide Batteries

Edited by Arthur Fleischer and J. J. Lander

Lead‐Acid Batteries

Hans Bode

Translated by R. J. Brodd and Karl V. Kordesch

Thin Films‐Interdiffusion and Reactions

Edited by J. M. Poate, M. N. Tu, and J. W. Mayer

Lithium Battery Technology

Edited by H. V. Venkatasetty

Quality and Reliability Methods for Primary Batteries

P. Bro and S. C. Levy

Techniques for Characterization of Electrodes and Electrochemical Processes

Edited by Ravi Varma and J. R. Selman

Electrochemical Oxygen Technology

Kim Kinoshita

Synthetic Diamond: Emerging CVD Science and Technology

Edited by Karl E. Spear and John P. Dismukes

Corrosion of Stainless Steels, Second Edition

A. John Sedriks

Semiconductor Wafer Bonding: Science and Technology

Q.‐Y. Tong and U. Göscle

Fundamentals of Electrochemistry, Second Edition

V. S. Bagotsky

Fundamentals of Electrochemical Deposition, Second Edition

Milan Paunovic and Mordechay Schlesinger

Uhlig's Corrosion Handbook, Third Edition

Edited by R. Winston Revie

Fuel Cells: Problems and Solutions

Vladimir S. Bagotsky

Lithium Batteries: Advanced Technologies and Applications

Edited by B. Scrosati, K. M. Abraham, W. A. van Schalkwijk, and J. Hassoun

Modern Electroplating, Fifth Edition

Edited by Mordechay Schlesinger and Milan Paunovic

Electrochemical Power Sources: Batteries, Fuel Cells, and Supercapacitors

By V. S. Bagotsky, A. M. Skundin, and Y. M. Volfkovic

Molecular Modeling of Corrosion Processes: Scientific Development and Engineering Applications

Edited by C. D. Taylor and P. Marcus

Atmospheric Corrosion, Second Edition

Christofer Leygraf, Inger Odnevall Wallinder, Johan Tidblad, and Thomas Graedel

Electrochemical Impedance Spectroscopy, Second Edition

Mark E. Orazem and Bernard Tribollet

Electrochemical Systems, Fourth Edition

John Newman and Nitash P. Balsara

ELECTROCHEMICAL SYSTEMS

Fourth Edition

JOHN S. NEWMAN and NITASH P. BALSARA

University of California, Berkeley

This edition first published 2021

© 2021 John Wiley & Sons Inc.

Edition History

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Library of Congress Cataloging‐in‐Publication Data

Hardback ISBN: 9781119514602

Central Cover Image: Redrawn from Figure 22.12 by Hee Jeung Oh, with permission from The Electrochemical Society

Cover Image: Courtesy of John Newman, Nitash P. Balsara, and Hee Jeung Oh

PREFACE TO THE FOURTH EDITION

Electrochemical systems provide the basis for many technologically important applications, such as batteries and fuel cells, production and refining of metals and chemicals, fabrication of electronic materials and devices, and operation of sensors, including those regulating the air/fuel ratio in automobile engines. The rechargeable lithium‐ion battery has emerged as a vital element of the emerging clean‐energy landscape. In biological systems, nerve action involves electrochemical processes. While applications continue to evolve, the fundamentals need only minor revision to train and guide people in adapting to new applications. Electrochemical systems involve many simultaneously interacting phenomena, drawn from many aspects of chemistry and physics, and require a disciplined learning process. The book provides a comprehensive coverage of electrochemical theories as they pertain to the understanding of electrochemical systems. It describes the foundations of thermodynamics, chemical kinetics, and transport phenomena including the electric potential and charged species.

This fourth edition incorporates further improvements developed over the years in teaching both graduate and advanced undergraduate students. Chapter 2 has expanded to include cells with polymer electrolytes. Chapter 6 now includes a discussion of equilibration of a charged polymer material and an electrolytic solution (Donnan equilibrium). The discussion of the oxygen electrode in Chapter 8 now includes insight from recent computer simulations. The application of concentrated solution theory to polymer electrolytes is added to Chapters 12 and 14. The number of transport properties describing different systems is now clearly stated. Chapter 15 presents a method for predicting turbulence by means of dissipation. Chapter 15 presents a method for predicting turbulence by means of dissipation. Finally, impedance measurements in electrochemical systems are important because experimental implementation is easy and diagnostic information is obtained without destroying the system. A new chapter on this subject, Chapter 24, is added.

We have much gratitude for the many students and colleagues who have done experiments and calculations that are reported in the book, and to our families for their continual support. We thank Saheli Chakraborty, Youngwoo Choo, Louise Frenck, Michael Galluzzo, Kevin Gao, Lorena Grundy, David Halat, Darby Hickson, Alec Ho, Zach Hoffman, Whitney Loo, Jacqueline Maslyn, Eric McShane, Hee Jeung Oh, Morgan Seidler, Gurmukh Sethi, Deep Shah, Neel Shah, and Irune Villaluenga, who patiently corrected many drafts of this manuscript. NPB thanks JN for the honor of working with him on the fourth edition and for being his mentor for more than a decade.

April 27, 2020

JOHN NEWMANBerkeley, California

NITASH P. BALSARABerkeley, California

PREFACE TO THE THIRD EDITION

This third edition incorporates various improvements developed over the years in teaching electrochemical engineering to both graduate and advanced undergraduate students. Chapter 1 has been entirely rewritten to include more explanations of basic concepts. Chapters 2, 7, 8, 13, 18, and 22 and Appendix C have been modified, to varying degrees, to improve clarity. Illustrative examples taken from real engineering problems have been added to Chapters 8 (kinetics of the hydrogen electrode), 18 (cathodic protection), and 22 (reaction‐zone model and flow‐through porous electrodes). Some concepts have been added to Chapters 2 (Pourbaix diagrams and the temperature dependence of the standard cell potential) and 13 (expanded treatment of the thermoelectric cell). The exponential growth of computational power over the past decade, which was made possible in part by advances in electrochemical technologies such as semiconductor processing and copper interconnects, has made numerical simulation of coupled nonlinear problems a routine tool of the electrochemical engineer. In realization of the importance of numerical simulation methods, their discussion in Appendix C has been expanded.

As discussed in the preface to the first edition, the science of electrochemistry is both fascinating and challenging because of the interaction among thermodynamic, kinetic, and transport effects. It is nearly impossible to discuss one concept without referring to its interaction with other concepts. We advise the reader to keep this in mind while reading the book, in order to develop facility with the basic principles as well as a more thorough understanding of the interactions and subtleties.

We have much gratitude for the many graduate students and colleagues who have worked on the examples cited and proofread chapters and for our families for their continual support. KET thanks JN for the honor of working with him on this third edition.

June 1, 2004

JOHN NEWMANBerkeley, California

KAREN E. THOMAS‐ALYEAManchester, Connecticut

PREFACE TO THE SECOND EDITION

A major theme of Electrochemical Systems is the simultaneous treatment of many complex, interacting phenomena. The wide acceptance and overall impact of the first edition have been gratifying, and most of its features have been retained in the second edition. New chapters have been added on porous electrodes and semiconductor electrodes. In addition, over 70 new problems are based on actual course examinations.

Immediately after the introduction in Chapter 1, some may prefer to study Chapter 11 on transport in dilute solutions and Chapter 12 on concentrated solutions before entering the complexities of Chapter 2. Chapter 6 provides a less intense, less rigorous approach to the potentials of cells at open circuit. Though the subjects found in Chapters 5, 9, 10, 13, 14, and 15 may not be covered formally in a one‐semester course, they provide breadth and a basis for future reference.

The concept of the electric potential is central to the understanding of the electrochemical systems. To aid in comprehension of the difference between the potential of a reference electrode immersed in the solution of interest and the electrostatic potential, the quasi‐electrostatic potential, or the cavity potential—since the composition dependence is quite different—Problems 6.16 and Figure 12.1 have been added to the new edition. The reader will also benefit by the understanding of the potential as it is used in semi‐conductor electrodes.

June 10, 1991

JOHN NEWMANBerkeley, California

PREFACE TO THE FIRST EDITION

Electrochemistry is involved to a significant extent in the present‐day industrial economy. Examples are found in primary and secondary batteries and fuel cells; in the production of chlorine, caustic soda, aluminum, and other chemicals; in electroplating, electromachining, and electrorefining; and in corrosion. In addition, electrolytic solutions are encountered in desalting water and in biology. The decreasing relative cost of electric power has stimulated a growing role for electrochemistry. The electrochemical industry in the United States amounts to 1.6 percent of all U.S. manufacturing and is about one third as large as the industrial chemicals industry.[1]

The goal of this book is to treat the behavior of electrochemical systems from a practical point of view. The approach is therefore macroscopic rather than microscopic or molecular. An encyclopedic treatment of many specific systems is, however, not attempted. Instead, the emphasis is placed on fundamentals, so as to provide a basis for the design of new systems or processes as they become economically important.