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Henry Taylor
Henry Taylor

Electrochemistry: Principles, Methods, and Applications, 2nd Edition



Electrochemistry 2nd Completely Revised and Updated Edition




Electrochemistry is a fascinating branch of chemistry that deals with the relationship between electrical energy and chemical changes. It has many applications in various fields, such as energy, environment, medicine, industry, and more. In this article, we will explore the basics of electrochemistry, its history and development, its principles and concepts, and its applications in different domains. We will also answer some frequently asked questions about electrochemistry.




Electrochemistry 2nd Completely Revised And Updated Edition


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What is electrochemistry?




Definition and scope of electrochemistry




Electrochemistry is the subdiscipline of chemistry that studies the chemical phenomena associated with charge separation, usually in liquid media, such as solutions. The separation of charge is often associated with charge transfer, which can occur homogeneously in solution between different chemical species, or heterogeneously on electrode surfaces.


Electrochemical reactions are those that involve the input or generation of electric currents. They are broadly classified into two categories:



  • Galvanic reactions: These are spontaneous reactions that produce electrical energy from chemical energy. For example, a battery is a device that converts the chemical energy stored in its electrodes into electrical energy that can power a circuit.



  • Electrolytic reactions: These are non-spontaneous reactions that require electrical energy to drive chemical changes. For example, electrolysis is a process that uses an external power source to split water into hydrogen and oxygen gas.



Electrochemistry covers a wide range of topics, such as thermodynamics, kinetics, transport phenomena, electrostatics, electrodynamics, materials science, surface science, analytical chemistry, and more. It also has many interdisciplinary connections with other fields, such as physics, biology, engineering, nanotechnology, and more.


History and development of electrochemistry




The history of electrochemistry can be traced back to ancient times, when people observed natural phenomena such as lightning, electric fish, static electricity, and magnetism. However, the scientific understanding of electrochemistry began in the 16th to 18th centuries, when scientists such as William Gilbert, Otto von Guericke, Stephen Gray, Benjamin Franklin, Luigi Galvani, Alessandro Volta, and others made discoveries and experiments related to electricity and magnetism.


In the 19th century, electrochemistry advanced significantly with the contributions of scientists such as Michael Faraday, Humphry Davy, Johann Wilhelm Ritter, Hans Christian Oersted, André-Marie Ampère, Georg Ohm, Gustav Kirchhoff, Robert Bunsen, William Grove, William Thomson, Rudolf Clausius, Walther Nernst, Svante Arrhenius, Jacobus van 't Hoff, Wilhelm Ostwald, and others. They established the laws and theories of electricity, magnetism, electrolysis, electromagnetism, electrochemical cells, thermodynamics, kinetics, equilibrium, and transport phenomena.


In the 20th century, electrochemistry continued to develop with the emergence of new fields and applications such as quantum mechanics, solid-state physics, electroanalytical chemistry, electrocatalysis, electroplating, corrosion science, fuel cells, batteries, solar cells, electrochemical sensors, bioelectrochemistry, and more. Some of the notable scientists who contributed to electrochemistry in this period include Niels Bohr, Albert Einstein, Erwin Schrödinger, Max Planck, Irving Langmuir, Frederick Donnan, Jaroslav Heyrovský, Charles Butler, John Bockris, Heinz Gerischer, Alan Bard, Allen Hill, and others.


Today, electrochemistry is a vibrant and active field that continues to explore new phenomena, concepts, methods, and applications. It also faces new challenges and opportunities in the context of global issues such as energy, environment, health, and sustainability.


Principles of electrochemistry




Oxidation and reduction reactions




The core of electrochemistry is the study of oxidation and reduction reactions, also known as redox reactions. These are reactions that involve a change in the oxidation state of one or more elements. The oxidation state is a measure of the degree of electron loss or gain by an atom in a compound.


When a substance loses an electron, its oxidation state increases; thus, it is oxidized. When a substance gains an electron, its oxidation state decreases; thus, it is reduced. For example, for the redox reaction


H2 + F2 2HF


can be rewritten as follows:


Oxidation reaction: H2 2H + 2e


Reduction reaction: F2 + 2e 2F


Overall reaction: H2 + F2 2H + 2F


In this case, H2 is being oxidized (and is the reducing agent or reductant), while F2 is being reduced (and is the oxidizing agent or oxidant). The following acronym is useful in remembering this concept: "OIL RIG"


"OIL RIG" is a useful mnemonic for remembering the definitions of oxidation and reduction.


Oxidation Is Losing electrons; Reduction Is Gaining electrons


Balancing redox reactions




To balance a redox reaction, we need to ensure that the number of atoms of each element and the total charge are conserved on both sides of the equation. There are different methods to balance redox reactions, such as the half-reaction method, the oxidation number method, and the ion-electron method. Here, we will illustrate the half-reaction method, which involves the following steps:



  • Separate the redox reaction into two half-reactions: one for oxidation and one for reduction.



  • Balance the atoms of each element (except oxygen and hydrogen) in each half-reaction.



  • Balance the oxygen atoms by adding H2O molecules to the side that needs oxygen.



  • Balance the hydrogen atoms by adding H ions to the side that needs hydrogen.



  • Balance the charge by adding electrons to the side that has a higher positive charge.



  • Multiply each half-reaction by an appropriate factor so that the number of electrons in both half-reactions are equal.



  • Add the two half-reactions together and cancel out any common terms on both sides.



  • If the reaction occurs in a basic medium, add OH ions to both sides to neutralize any H ions. Then combine H and OH to form H2O and cancel out any common terms on both sides.



  • Simplify the final equation if possible.



To illustrate this method, let us balance the following redox reaction that occurs in an acidic medium:


MnO4(aq) + C2O4(aq) Mn(aq) + CO2(g)


The steps are as follows:



  • The two half-reactions are: sup>+2(aq) C2O4(aq) CO2(g)



  • The atoms of each element (except oxygen and hydrogen) are already balanced in each half-reaction.



  • To balance the oxygen atoms, we add H2O molecules to the side that needs oxygen: MnO4(aq) Mn(aq) + 4H2O(l) C2O4(aq) 2CO2(g) + 2H2O(l)



  • To balance the hydrogen atoms, we add H ions to the side that needs hydrogen: MnO4(aq) + 8H(aq) Mn(aq) + 4H2O(l) C2O4(aq) 2CO2(g) + 2H2O(l) + 2H(aq)



  • To balance the charge, we add electrons to the side that has a higher positive charge: MnO4(aq) + 8H(aq) + 5e Mn(aq) + 4H2O(l) C2O4(aq) 2CO2(g) + 2H2O(l) + 2H(aq) + 2e



  • To make the number of electrons equal in both half-reactions, we multiply the first half-reaction by 2 and the second half-reaction by 5: 2[MnO4(aq) + 8H(aq) + 5e Mn(aq) + 4H2O(l)] 5[C2O4(aq) 2CO2(g) + 2H2O(l) + 2H(aq) + 2e]



  • We add the two half-reactions together and cancel out any common terms on both sides: (aq) + 2MnO4(aq) 10CO2(g) + 8H2O(l) + 2Mn(aq)



  • Since the reaction occurs in an acidic medium, we do not need to add OH ions to both sides.



  • The final equation is already simplified.



The balanced redox reaction is:


10e + 16H(aq) + 5C2O4(aq) + 2MnO4(aq) 10CO2(g) + 8H2O(l) + 2Mn(aq)


Electrochemical cells




An electrochemical cell is a device that converts chemical energy into electrical energy or vice versa by means of a redox reaction. There are two types of electrochemical cells: galvanic cells and electrolytic cells.


A galvanic cell is a type of electrochemical cell that produces electrical energy from a spontaneous redox reaction. It consists of two electrodes (metallic conductors) immersed in an electrolyte (a solution that contains ions). The electrodes are connected by an external circuit that allows the flow of electrons. The electrodes are also connected by a salt bridge or a porous membrane that allows the flow of ions and maintains the electrical neutrality of the solutions.


In a galvanic cell, the electrode where oxidation occurs is called the anode and the electrode where reduction occurs is called the cathode. The anode has a negative charge and the cathode has a positive charge. The electrons flow from the anode to the cathode through the external circuit, generating an electric current. The ions flow from one solution to another through the salt bridge or the porous membrane, completing the circuit.


A common example of a galvanic cell is the Daniell cell, which consists of a zinc electrode immersed in a solution of zinc sulfate and a copper electrode immersed in a solution of copper sulfate. The redox reaction that occurs in this cell is:


Zn(s) + Cu(aq) Zn(aq) + Cu(s)


This reaction can be separated into two half-reactions:


Anode (oxidation): Zn(s) Zn(aq) + 2e


Cathode (reduction): Cu(aq) + 2e Cu(s)


The schematic diagram of a Daniell cell is shown below:


![Daniell cell](https://upload.wikimedia.org/wikipedia/commons/thumb/9/9b/Daniell_cell_%28US%29.svg/1200px-Daniell_cell_%28US%29.svg.png) An electrolytic cell is a type of electrochemical cell that consumes electrical energy to drive a non-spontaneous redox reaction. It also consists of two electrodes immersed in an electrolyte and connected by an external circuit. However, in this case, the electrodes are connected to an external power source that provides the electric current. The electrodes are also labeled differently: the electrode connected to the positive terminal of the power source is called the anode and the electrode connected to the negative terminal of the power source is called the cathode.


the external circuit. The ions flow from one solution to another through the salt bridge or the porous membrane, completing the circuit.


A common example of an electrolytic cell is the electrolysis of water, which consists of two inert electrodes (such as platinum) immersed in a solution of water and a small amount of an electrolyte (such as sulfuric acid) to increase the conductivity. The redox reaction that occurs in this cell is:


2H2O(l) 2H2(g) + O2(g)


This reaction can be separated into two half-reactions:


Anode (oxidation): 2H2O(l) O2(g) + 4H(aq) + 4e


Cathode (reduction): 4H(aq) + 4e 2H2(g)


The schematic diagram of an electrolytic cell for water electrolysis is shown below:


![Electrolytic cell for water electrolysis](https://upload.wikimedia.org/wikipedia/commons/thumb/0/0a/Electrolysis_of_Water.svg/1200px-Electrolysis_of_Water.svg.png) Standard electrode potential




The standard electrode potential (E) is a measure of the tendency of a half-reaction to occur as a reduction reaction at standard conditions (25C, 1 atm, and 1 M concentrations). It is expressed in volts (V) and is determined by measuring the voltage of a galvanic cell that consists of a standard hydrogen electrode (SHE) as the reference electrode and the electrode of interest as the working electrode.


The SHE is an electrode that consists of a platinum wire coated with platinum black and immersed in a solution of 1 M H. It is connected to a hydrogen gas inlet at 1 atm pressure. The SHE is assigned a potential of 0 V by convention. The working electrode is an electrode that consists of a metal or a metal ion that participates in the half-reaction of interest. The working electrode is immersed in a solution of 1 M metal ion.


The standard cell potential (Ecell) is the difference between the standard electrode potentials of the cathode and the anode. It is calculated by using the following equation:


Ecell = Ecathode - Eanode


The standard cell potential can be used to determine the spontaneity of a redox reaction. If Ecell is positive, then the reaction is spontaneous. If Ecell is negative, then the reaction is non-spontaneous. If Ecell is zero, then the reaction is at equilibrium.


A common example of a galvanic cell that uses standard electrode potentials is the standard hydrogen-zinc cell, which consists of a SHE as the cathode and a zinc electrode as the anode. The redox reaction that occurs in this cell is:


Zn(s) + 2H(aq) Zn(aq) + H2(g)


This reaction can be separated into two half-reactions:


H2(g) Ecathode = 0 V (by definition)


Anode (oxidation): Zn(s) Zn(aq) + 2e Eanode = -0.76 V (from a table of standard electrode potentials)


The standard cell potential is:


Ecell = Ecathode - Eanode


Ecell = 0 V - (-0.76 V)


Ecell = 0.76 V


Since Ecell is positive, the reaction is spontaneous.


Spontaneity of redox reactions




The spontaneity of a redox reaction can be predicted by using the standard cell potential, as explained above. However, the standard cell potential only applies to standard conditions. In reality, the conditions may vary from the standard ones, such as the temperature, pressure, and concentrations of the reactants and products. Therefore, we need to use another criterion to determine the spontaneity of a redox reaction under non-standard conditions. This criterion is the Gibbs free energy change (ΔG).


The Gibbs free energy change is a measure of the amount of energy available to do work in a system. It is related to the standard cell potential and the reaction quotient (Q) by the following equation:


ΔG = ΔG + RTlnQ


where ΔG is the standard free energy change, R is the universal gas constant, T is the absolute temperature, and Q is the reaction quotient.


The reaction quotient is a ratio of the concentrations or partial pressures of the products and reactants at any given time. It is calculated by using the following expression:


Q = [C][D]/[A][B]


where [A], [B], [C], and [D] are the concentrations or partial pressures of the reactants and products, and a, b, c, and d are their stoichiometric coefficients.


The Gibbs free energy change can be used to determine the spontaneity of a redox reaction under non-standard conditions. If ΔG is negative, then the reaction is spontaneous. If ΔG is positive, then the reaction is non-spontaneous. If ΔG is zero, then the reaction is at equilibrium.


A common example of a redox reaction that occurs under non-standard conditions is the corrosion of iron in air and water. The redox reaction that occurs in this process is:


4Fe(s) + 3O2(g) + 6H2O(l) 4Fe(OH)3(s)


This reaction can be separated into two half-reactions:


(aq) + 4e 2H2O(l) Ecathode = 1.23 V (from a table of standard electrode potentials)


Anode (oxidation): Fe(s) Fe(aq) + 2e Eanode = -0.44 V (from a table of standard electrode potentials)


The standard cell potential is:


Ecell = Ecathode - Eanode


Ecell = 1.23 V - (-0.44 V)


Ecell = 1.67 V


The standard free energy change is:


ΔG = -nFEcell


ΔG = -(8 mol)(96.5 kJ/mol)(1.67 V)


ΔG = -1293 kJ/mol


The reaction quotient is:


Q = [Fe]/[O2]


The free energy change under non-standard conditions is:


ΔG = ΔG + RTlnQ


)


Depending on the values of [Fe] and [O2], ΔG may be negative, positive, or zero. However, in most cases, ΔG is negative, meaning that the reaction is spontaneous and iron corrodes in air and water.


Cell emf dependency on changes in concentration




The cell emf (electromotive force) is the voltage or potential difference between the two electrodes of an electrochemical cell. It is related to the standard cell potential and the reaction quotient by the Nernst equation:


Ecell = Ecell - (RT/nF)lnQ


where Ecell is the standard cell potential, R is the universal gas constant, T is the absolute temperature, n is the number of moles of electrons transferred in the reaction, F is the Faraday constant, and Q is the reaction quotient.


The Nernst equation shows how the cell emf depends on the changes in concentration of the reactants and products. As Q increases, lnQ increases, and Ecell decreases. As Q decreases, lnQ decreases, and Ecell increases. When Q equals 1, lnQ equals 0, and Ecell equals Ecell.


A common example of a concentration cell is a galvanic cell that consists of two electrodes of the same metal but immersed in solutions of different concentrations of the same metal ion. For example, a copper concentration cell consists of two copper electrodes immersed in solutions of 1 M and 0.1 M Cu, respectively. The redox reaction that occurs in this cell is:


Cu(s) + Cu(aq) Cu(s) + Cu(aq)


This reaction can be separated into two half-reactions:


Anode (oxidation): Cu(s) Cu(aq) + 2e


Cathode (reduction): Cu(aq) + 2e Cu(s)


The standard cell potential is:


Ecell = Ecathode - Eanode


Ecell = 0 V - 0 V (since both electrodes have the same standard electrode potential)


the environment. Some examples are stainless steel crevice corrosion in chloride solutions, aluminum crevice corrosion in seawater, and copper crevice corrosion in acidic water.


  • Stress corrosion cracking: This is a type of catastrophic corrosion that occurs when the metal or the alloy is exposed to an environment that causes cracks to form and propagate on its surface. The cracks act as stress concentrators and corrode faster than the surrounding metal or alloy. The cracks may cause sudden and unexpected failure of the metal or the alloy. The corrosion rate depends on the applied or residual stress, the composition and structure of the metal or alloy, and the nature and concentration of the corrosive agent. Some examples are steel cracking in ammonia solutions, brass cracking in ammonia solutions, and titanium cracking in methanol solutions.



Corrosion can be prevented or controlled by various methods, such as coating, cathodic protection, sacrificial anodes, corrosion inhibitors, material select


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