Technology & Engineering
Anodic Cathodic Reaction
Anodic and cathodic reactions are fundamental processes in electrochemistry. The anodic reaction involves the loss of electrons, leading to oxidation, while the cathodic reaction involves the gain of electrons, leading to reduction. These reactions occur at separate electrodes in an electrochemical cell and are essential for processes such as corrosion, batteries, and electrolysis.
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9 Key excerpts on "Anodic Cathodic Reaction"
- S Borenstein(Author)
- 1994(Publication Date)
- Woodhead Publishing(Publisher)
0 A chemical reaction: a reaction that does not involve the transfer of electrons. An electrochemica1 reaction: a chemical reaction involving the transfer of electrons. 113 114 Microbiologically influenced corrosion handbook 0 0 0 0 0 0 0 0 0 0 0 0 0 0 Electrode: a metal in contact with an electrolyte at the location where the current either enters or leaves the metal to enter the solution. Electrolyte: a medium that conducts electric current by ionic movement. Anions: negatively charged ions attracted to the anode. Cations: positively charged ions attracted to the cathode. Anode: the portion of the metal where current is discharged and corrosion occurs. Cathode: the metal surface where electrons from the anode (by way of the electronic path) react with the electrolyte. The pH: measurement of the hydrogen ion concentration. It gen- erally denotes the degree of acidity or basicity of a solution: Oxidation: this process occurs when electrons are lost from an atom or compound. Oxidation decreases the negative charge (or increases the positive charge). Oxidation occurs at the anode. Reduction: this process occurs when electrons are added to an atom or compound. Reduction increases the negative charge (or decreases the positive charge). Reduction occurs at the cathode. Half cell reactions: two separate reactions, anodic and cathodic, occurring simultaneously, owing to the immersion of an electrode in an electrolyte designed to measure a single electrode’s potential. Polarization: the flow of potential because of current flow. Concentration polarization: the retarding of an electrochemical reaction as a result of concentration changes in the solution adjacent to the metal surface. Cathodic polarization: when a cathodic reaction product such as, for example, hydrogen, is not removed. Activation polarization: the retarding actions that are inherent to the reaction itself.
eBook - PDF- G R Jones(Author)
- 2013(Publication Date)
- Newnes(Publisher)
In electroche-mical reactors the reverse is true, i.e. anodic reaction at the positive and cathodic at the negative electrode. 40.2.1 Redox process The chemical reactions at the electrodes are either 'reduction' or 'oxidation', i.e. 'redox' processes. The basic feature of such reactions is the gain or loss, respectively, of one or more electrons, e.g. reduction/cathodic reaction 2H + + 2e~ -» H 2 Cu 2+ + 2e~ -> Cu oxidation/anodic reaction Zn —> Zn 2+ + 2e~ Pb -> Pb 2+ + 2e~ Here e represents the electron. H, Cu, Zn and Pb represent atoms, or (when charged) ions of hydrogen, copper, zinc and lead. The sign + indicates a deficiency of one electron; 2+ indicates a deficiency of two electrons. The gaseous hydrogen formed by the first reduction pro-cess, if allowed to accumulate, would rapidly polarise the electrode, and the electrochemical reaction would virtually cease. To overcome this, the positive electrodes of many kinds of batteries are selected from substances which readily under-go a depolarising reaction with hydrogen. Typical examples are manganese dioxide (Mn0 2 ), used in the primary dry battery, and lead dioxide (Pb0 2 ), used in the leacy acid storage battery. Negative electrodes must be readily oxidisable, and for these active metals such as zinc, lead and cadmium are generally chosen. Electrode potentials and redox reactions are not confined to metals, but include such elements as hydrogen, oxygen, chlorine or fluorine, either in the gaseous form or in combina-tion with some metal in the form of an inorganic salt. Electrode potentials are generally expressed with respect to hydrogen, which for standard conditions is assigned a poten-tial of zero. It is an advantage to have at the two electrodes redox reactions widely spaced on the potential scale to give the highest cell e.m.f. But the choice is restricted by other factors—in particular, the type of electrolyte. Formerly elec-trolytes were aqueous solutions of salts, acids or bases.
eBook - PDF- D Pletcher, R Greff, R Peat, L M Peter, J Robinson(Authors)
- 2001(Publication Date)
- Woodhead Publishing(Publisher)
1.1 — Schematic view of some types of electrode reactions met in applied and fundamental electrochemistry. 18 Introduction to the fundamental concepts of electrochemistry [Ch. 1 Electrolysis is only possible in a cell with both an anode and a cathode, and, because of the need to maintain an overall charge balance, the amount of reduction at the cathode and oxidation at the anode must be equal. The total chemical change is found by adding the two individual electrode reactions; for example, the chemical change in a chlor-alkali membrane or diaphragm cell is obtained by adding Equations (1.3) and (1.8), i.e. 2CP+ 2H 2 0 -> Cl 2 + H 2 + 20H~ . (1.11) Moreover, when electrolysis occurs, in addition to electron transfer at the anode and cathode surfaces, ions must pass through the solution between the electrodes and electrons though the wires externally interconnecting the two electrodes (in order to maintain electrical neutrality at all points in the system). Hence the current through the external circuit, /, given by / = AI (1.12) (when A is the electrode area and / the current density), is a convenient measure of the rate of the cell reaction, and the charge, q, passed during a period, f, indicates the total amount of chemical reaction which can have taken place: indeed, the charge required to convert m moles of starting material to product in an electrode reaction involving the transfer of n electrons/molecule is readily calculated using Faraday's law rt q = I idt = mnF . (1.13) Jo When the two electrodes of a cell are interconnected by an external circuit, however, the cell reaction will only occur spontaneously if the free energy change associated with the net cell reaction is negative. This is not the case in a cell for the production of chlorine and caustic soda, i.e. the free energy of reaction (1.11) is positive, and for reaction (1.11) to occur it will be necessary to supply energy by applying a potential between the two electrodes. This potential must certainly be
eBook - PDF- Keith Oldham, Jan Myland(Authors)
- 2012(Publication Date)
- Academic Press(Publisher)
Likewise, the electroreactant at a cathode This possibly dates back to Faraday who wrote in 1834 that ... anions which go to the anode ... and those passing to the cathode, cations ... . However, Faraday meant something different by anions and cations from what we mean today. 5 Electrode Reactions may be any reducible species: cation, neutral, or anion, as exemplified by the reactions 5:2:4 Fe 3 + (aq) + e » Fe 2 + (aq) 5:2:5 H 2 0 2 (aq) + 2 e » 20H(aq) and 5:2:6 I 3 (aq) + 2 e > 31 (aq) all of which occur at inert electrodes in contact with aqueous solutions. In the above six examples of electrode reactions, the product, like the reactant, is dissolved in the ionic conductor (aqueous solutions in these cases). However, instances occur in which the electroproduct dissolves in the electronic conductor. Examples were encountered in our discussions of the chloralkalki industry (page 94) and of anodic stripping (page 134) in which metals, produced by electrode reactions, dissolve in the mercury cathode to form amalgams 5:2:7 M n + (aq) + «e~(Hg) > M(amal) A different example is provided by the electrolysis of aqueous solutions using a palladium cathode. The hydrogen produced remains as atoms and enters the lattice of the palladium cathode 5:2:8 H 2 0(t) + e(Pd) > H(Pd) + OH(aq) Within the lattice each hydrogen atom ionizes, the electron joining those of the palladium metal, and the resulting proton H + is free to move. Instead of dissolving in one or other of the phases that bound the electrode interface, the electroproduct may form a new phase. For example, a metal such as silver may be cathodically electrodeposited onto other metals by the reaction 5:2:9 Ag(CN)~(aq) + e » Ag(s) + 2CN(aq) This is an example of an electrode reaction used in electroplating (Section 3:7). Frequently the new phase is a gas, as with 5:2 Types of electrode reaction 153 5:2:10 2Cl(aq) > 2e~ + Cl 2 (g) the electrooxidation which occurs at carbon anodes in the chloralkali industry.
eBook - PDF- Thomas F. Fuller, John N. Harb(Authors)
- 2018(Publication Date)
- Wiley(Publisher)
Hence, this chapter will introduce you to the key vocabu- lary of the discipline, as well as to some of the central aspects of electrochemical systems. In order to do this, we begin by looking at an electrochemical cell. 1.1 ELECTROCHEMICAL CELLS Electrochemical cells, such as the cell illustrated in Figure 1.1, lie at the heart of electrochemical systems. A typical electrochemical cell consists of two electrodes: an anode where oxidation occurs and a cathode where reduction takes place. Electrons move through an external circuit via an electronic conductor that connects the anode and cathode. The liquid solution that is between the two electrodes is the electrolyte. The electrolyte does not conduct electrons and does not contain any free electrons. It does, however, contain a mixture of negatively charged ions (anions) and positively charged ions (cations). These ions are free to move, which allows them to carry current in the electrolyte. The reactions take place at the electrode surface and are called heterogeneous electron-transfer reactions. For example, the electrodeposition of copper in the cell shown in Figure 1.1 can be written as Cu 2 + aq + 2e Cu s (1.1) Copper ions in solution accept two electrons from the metal and form solid copper. The reaction is described as heterogeneous because it takes place at the electrode surface rather than in the bulk solution; remember, there are no free electrons in the solution. Importantly, then, we see that electrochemical reactions are surface reactions. The metal that accepts or supplies electrons is the elec- trode. As written, copper ions gain electrons and therefore are reduced to form copper metal. When reduction occurs, the electrode is called the cathode. In the same cell, the reaction that takes place on the other electrode is Zn s Zn 2 + aq + 2e (1.2) The zinc metal is oxidized, giving up two electrons and forming zinc ions in solution. When oxidation occurs at the surface, the electrode is called the anode.
eBook - PDF- Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
As a result, the anode carries a slight negative charge and the cathode a slight positive charge. • In an electrolytic cell, the situation is reversed. Here, oxidation at the anode is forced to occur, which requires that the anode is positive so it can remove electrons from the reactant at that electrode. On the other hand, the cathode must be made negative so it can force the reactant at the electrode to accept electrons. By agreement among scientists, the names anode and cathode are always assigned according to the nature of the reaction taking place at the electrode. • If the reaction is oxidation, the electrode is called the anode. • If the reaction is reduction, the electrode is called the cathode. It is important to remember: • in an electrolytic cell, the cathode is negative (reduction) and the anode is positive (oxidation) • in a galvanic cell, the cathode is positive (reduction) and the anode is negative (oxidation). Electrolysis in aqueous solutions When electrolysis is carried out in an aqueous solution, the electrode reactions can be more complicated; we must consider oxidation and reduction of the solute as well as oxidation and reduction of water. (It is important to note that the nature of the electrode material itself can also strongly influence the outcome of the electrolysis. We will not discuss this further here.) For example, electrolysis of a solution of potassium sulfate (figure 12.25) gives hydrogen and oxygen. At the cathode, water is reduced, not K + . 2H 2 O(l) + 2e - → H 2 (g) + 2OH - (aq) (cathode) At the anode, water is oxidised, not the sulfate ion. 2H 2 O(l) → O 2 (g) + 4H + (aq) + 4e - (anode) CHAPTER 12 Oxidation and reduction 607 FIGURE 12.25 Electrolysis of an aqueous solution of potassium sulfate, K 2 SO 4 . The products of the electrolysis are the gases hydrogen and oxygen.
eBook - PDF- Dieter Landolt(Author)
- 2007(Publication Date)
- EPFL PRESS(Publisher)
We shall look here at the kinetics of the reduction of protons, which determines the corrosion rate in acid environments (Chapter 4). The cathodic reaction involves two electrons: 2 H + 2 H + 2 e → (5.39) The charge-transfer reaction at the metal-electrolyte interface proceeds in two steps. First a proton is reduced to an adsorbed hydrogen atom (Volmer reaction). In a second step, a hydrogen molecule forms, either by chemical reaction between two adsorbed hydrogen atoms (Tafel reaction) or via an electrochemical reaction between an adsorbed hydrogen atom and a proton (Heyrovsky reaction). Finally, the hydrogen molecule leaves the cathode, either in the form of gas bubbles or by diffusion in the electrolyte. Under certain conditions, the adsorbed hydrogen atoms can dissolve into the metal instead of forming hydrogen molecules (See Chapter 11). Volmer-Heyrovsky Mechanism According to this mechanism, the formation of a hydrogen molecule proceeds in two consecutive electrochemical steps. step I (RDS): H + H + ads e → (5.40) step II: H + H + H + ads 2 e → (5.41) (I) + (II): 2 H + 2 H + 2 e → (5.42) Step (II) is a charge-transfer reaction, in which the adsorbed hydrogen atom reacts with a proton and with an electron provided by the electrode. If step (I), called Volmer reaction, is rate limiting and, if in addition, α = α = α and α = α = α , the I,a II,a a I,c II,c c reaction rate, v , can be expressed as follows: I v v v k f E k c I I, a I, c a,I H a c,I H H + = − = ( ) − − ( ) θ α θ exp e 1 xp − ( ) α c f E (5.43) The rate of the anodic partial reaction, v , is proportional to the amount of adsorbed I,a hydrogen, expressed by the coverage θ , which represents the fraction of adsorption H sites occupied by hydrogen (0 < θ < 1). The rate of the cathodic partial reaction H v is proportional to the fraction of non-occupied sites, 1–θ . The cathodic partial I,c H current density thus becomes:
eBook - PDF- Yongchang Huang, Jianqi Zhang, Yongchang Huang, Jianqi Zhang(Authors)
- 2018(Publication Date)
- De Gruyter(Publisher)
Yongchang Huang 3 Electrochemical corrosion kinetics 3.1 Corrosion potential and its polarization diagram 3.1.1 Corrosion potential When metals suffer from electrochemical corrosion in aqueous solutions, their surfaces have at least two different electrode reactions at the same time, one of which is the metal oxidation reaction (anodic process) and another which is the reduction reaction of the oxidants in the solution (cathodic process). The rates of these two reac- tions are equal so that there is no accumulation of charges on the metal surface and the net current is zero. The potential of the corroded metal stays constant after metal immersion in an agent for a long time. The absolute potential cannot be measured currently, while its relative value can be measured by the reference electrode. The electromotive force of the galvanic cell constituted by the reference electrode and the electrode to be measured is the potential of the electrode to be measured. Now, let us continue to discuss the meaning of the electrode potential when there are two electrode reactions on a metal at the same time. As the net current on the metal is zero, it can be recognized as an isolate electrode. Below, we will analyze the problem by an imaginary short-circuited primary cell. In Fig. 3.1(a), the anode and cathode are made up by the same metal M, and the two electrodes are linked by a wire whose resistance is nearly zero, which means the anodic reaction and cathodic reaction are likely to occur on one metal surface, shown in Fig. 3.1(b). Fig. 3.1: Short-circuited primary cells constituted by the same metal [1]. (a) Primary cell whose electrode materials are made up by one metal. (b) Relationship between the primary cell constituted by the electrodes made up by same metal and the coupled system of the electrode reactions. https://doi.org/10.1515/9783110310054-003
eBook - PDFBatteries for Electric Vehicles
Materials and Electrochemistry
- Helena Berg(Author)
- 2015(Publication Date)
- Cambridge University Press(Publisher)
I Electrochemistry and battery technologies 1 The electrochemical cell The most fundamental unit of a battery is the electrochemical cell. All performance characteristics are dependent on the materials inside the cell, and all cells work according to some general principles independent of the materials employed. The purpose of this chapter is to bring together the fundamental aspects of an electrochemical cell as the basis for all further steps in the development of a battery intended for electric vehicles. An electrochemical cell converts chemical energy to electric energy when discharged, and vice versa. In addition, the electrochemical cells can be said to be either electrolytic or galvanic. In an electrolytic cell, the electric energy is converted to chemical energy (charging of the battery) and in a galvanic cell chemical energy is converted to electric energy (discharging of the battery). The basic design of an electrochemical cell consists of a positive and a negative electrode separated by an electrolyte, as shown in Figure 1.1. The chemical reactions taking place during charge and discharge processes are based on electrochemical oxidation and reduction reactions, known as the redox reactions, at the two electrodes. In these reactions, electrons are transferred via an external circuit from one electrode to another, and at the same time ions are transferred inside the cell, through the electrolyte, to maintain the charge balance. The species oxidised is called the oxidant, and the species reduced is called the reductant. The oxidation reaction takes place at the negative electrode, the anode, and electrons are transferred, via the external circuit, to the positive electrode, the cathode, where the reduction reaction takes place by accepting the electrons. The negative electrode is thus an electron donor, and the positive electrode an electron acceptor. During charge and discharge of a battery, the nomenclature of the electrodes changes.
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