Chemistry
Galvanic and Electrolytic Cells
Galvanic and electrolytic cells are two types of electrochemical cells. In a galvanic cell, spontaneous chemical reactions produce electrical energy, while in an electrolytic cell, electrical energy is used to drive non-spontaneous chemical reactions. Galvanic cells are commonly used in batteries to provide power, while electrolytic cells are used in processes like electroplating and electrolysis.
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10 Key excerpts on "Galvanic and Electrolytic Cells"
eBook - PDF- Ilana Fried(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
2. The Galvanic Cell, Basic Definitions and Concepts Every galvanic and electrolytic cell has two electrodes. The one which introduces electrons to the cell is the cathode; the one which removes them from the cell is the anode. In this chapter we will consider some of the basic concepts of cells; the meaning of cell and electrode potentials will be examined in detail. We shall define electrolytic cells and galvanic cells; discuss the difference between surfaces and bulk of matter; consider the location of the site of the electrode reaction, and the forces and laws which control the flow of current and make one electrode the source and the other the sink for electrons. A. The cell A galvanic cell is a device which generates electrical energy directly from chemical reactions. An electrolytic cell is a device by which a chemical reaction is made to take place by using electrical energy. These two devices are the same instrument operated in opposite ways and will be called here a cell. Figure 1 shows two cells connected in series. One operates as a galvanic cell and the other as an electrolytic cell. These are made of suitable containers which are each divided into two halves by a porous membrane. In the cell marked G the two halves contain solutions of copper and zinc sulphate with copper and zinc metal rods dipping into their respective solutions. The electrode reactions are cathode C u 2 + + 2 e = Cu I anode Zn = Z n 2 + + 2 e II If the salt solution is fairly concentrated, say of one molar, a potential difference of approximately 1-1 volts will be produced between the metals. In the cell marked Ε there are platinum electrodes dipping into solutions which do not contain platinum ions. The right-hand compartment contains ferric and ferrous ions and the electrode reaction is F e 3 + + e = F e 2 + III The platinum acts here merely as an electron membrane, its potential being determined by the concentration ratio of ferric and ferrous ions in the 9
eBook - PDFFuel Cells
Principles, Design, and Analysis
- Shripad T. Revankar, Pradip Majumdar(Authors)
- 2016(Publication Date)
- CRC Press(Publisher)
35 Review of Electrochemistry Sir Humphry Davy (1778–1829) Michael Faraday (1791–1867) Sir William Robert Grove (1811–1896) Svante August Arrhenius (1859–1927) Walther Hermann Nernst (1864–1941) 36 Fuel Cells 2.1 Electrochemical and Electrolysis Cell An electrochemical cell involves a chemical reaction driven by electricity or the cell generates electricity because of spontaneous chemical reaction (Bockris and Reddy, 1970; Bogotsky, 2006). The cell that produces electricity because of spontaneous chemical reaction is called a voltaic cell or a galvanic cell . The cell in which chemical reaction is nonspontaneous and is forced by external electricity is called an electrolysis cell . In the galvanic cell, the current is caused by the reactions releasing and accepting electrons at the different ends of a conductor. A common example of a galvanic cell is a bat-tery. In contrast to this electrolysis, cells decompose chemical compounds by means of electrical energy through electrolysis with a net increase in chemical energy. An example of electrolysis is the decomposition of water into hydrogen and oxygen gas with external voltage supply. The galvanic cell consists of two separate compartments called half-cells. The two half-cells may use the same electrolyte, or they may use different electrolytes. Each half-cell consists of electrolyte solutions and electrodes that can be connected in a circuit to some voltmeter placed between the two electrodes within the circuit. In one half-cell, an electrode called the anode accepts electrons given by ion species (anions) that migrate to the electrode that are then passed through conducting wires in a circuit. In the other half-cell, the electrode called the cathode attracts ions (cations) where electrons can be gained by the species that migrates to that electrode. In order for this to occur, the circuit is completed by a conducting medium that allows ions or electrons to pass from one half-cell to the other.
eBook - ePubElectrical Engineering
Fundamentals
- Viktor Hacker, Christof Sumereder(Authors)
- 2020(Publication Date)
- De Gruyter Oldenbourg(Publisher)
6 Electrochemistry6.1 Basic electrochemical concepts
With special regard to electrical engineering, this chapter covers the branch of electrochemistry that deals with the generation and storage of electric current. The electrochemical oxidation and reduction reactions take place at the phase boundaries of the electrode and the electrolyte.Galvanic cellChemical energy is transformed into electrical energy, current is produced, and electrochemical reactions take place spontaneously (negative free enthalpy). Galvanic cells are categorised into three subgroups:- Primary cells
- Secondary cells
- Fuel cells
Electrolytic cellElectric energy is transformed into chemical energy. Two electrodes made of electron-conducting material, and the electrolytes with ion conductivity are conductively connected62 to each other. At the two spatially separated electrodes electrochemical reactions take place.Half-cellA half-cell consists of one single electrode and an electrolyte into which the electrode is submerged (e.g. copper in a copper sulphate solution). If a (metal) electrode is submerged into a metal salt solution (same metal), the surface of the electrode becomes charged. With base metals (e.g. zinc) some metal atoms enter the solution and the released electrons stay on the surface of the electrode, which is now negatively charged. The positively charged metal ions remain bound to the negatively charged metal surface. Thereby an electrical double layer is formed where the negative and the positive charges balance each other out. When two half-cells are combined, a galvanic cell (connected through ionic conductor and electron conductor) is formed.Anode
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2014(Publication Date)
- Wiley(Publisher)
Electrochemistry Chapter Outline 19.1 | Galvanic (Voltaic) Cells 19.2 | Cell Potentials 19.3 | Utilizing Standard Reduction Potentials 19.4 | E ° cell and ∆G° 19.5 | Cell Potentials and Concentrations 19.6 | Electricity 19.7 | Electrolytic Cells 19.8 | Electrolysis Stoichiometry 19.9 | Practical Applications of Electrolysis People around the world are coming to the conclusion that we must develop “green” economic and consumption policies that are sustainable far into the future. This means that the use of resources should be minimized, and those resources that are used for a wide variety of human activities must be replenished rapidly. One way to achieve a sustainable energy economy is to convert solar and wind energy directly into electricity, using technology such as the windmills shown in the chapter opening photograph, bypassing petroleum and all of its production, refining, and pollution problems. Since solar energy is only available for a part of a day and wind energy is variable, economical storage of electrical energy in high-tech batteries is essential. 19 904 Erik Isakson/Getty Images 19.1 | Galvanic (Voltaic) Cells 905 19.1 | Galvanic (Voltaic) Cells 905 19.1 | Galvanic (Voltaic) Cells Batteries have become common sources of portable power for a wide range of consumer products, from cell phones to iPods to laptops and hybrid cars. The energy from a battery comes from a spontaneous redox reaction in which the electron transfer is forced to take place through a wire. The apparatus that provides electricity in this way is called a galvanic cell, after Luigi Galvani (1737–1798), an Italian anatomist who discovered that electricity can cause the contraction of muscles.
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.
eBook - PDFPhysical Chemistry
Thermodynamics
- Horia Metiu(Author)
- 2006(Publication Date)
- Taylor & Francis(Publisher)
28 GALVANIC CELLS: PHENOMENA Introduction §1. In this and the next chapter, you will learn some electrochemistry. This is the branch of chemistry in which chemical reactions are used to produce an electrical current or an electrical current is used to cause chemical reactions. This is an old field, started by Galvani, Volta, Davy, and Faraday two centuries ago. Electrochemistry is used in many industrial processes: the fabrication of aluminum, chlorine, NaOH, and many specialty chemicals. It is also useful to analytical chemistry. The recent growth in the number of portable electronic devices has increased the demand for better batteries. Much research is being done for producing batteries that are lighter, last longer, and require less time for recharging. Of particular inter-est are fuel cells, which function by oxidizing a variety of fuels (e.g., H 2 , methanol) to directly produce electricity. These batteries do not need recharging, but only a continuous supply of fuel. Some people anticipate that in the next few decades, fuel cells will propel cars and other mobile, heavy equipment, run solid state electronic devices, and function as sources of electricity. This will be true if the research done within the next two decades is successful. Since this is not a book on electrochemistry, I can only give the subject two chapters. One describes electrochemical cells and their function, and the other applies thermodynamics to establish a connection between the maximum voltage 571 572 Galvanic Cells: Phenomena produced by a cell and its chemical nature and electrolyte concentration. This has unexpected consequences: electrochemical cells are the best devices for studying chemical equilibrium in redox reactions and for measuring the activity coefficients of electrolyte solutions. Galvanic Cells §2. How to Make a Daniell Cell. To introduce you to electrochemical cells, let us make one.
eBook - PDF- Krystyna Jackowska, Paweł Krysiński(Authors)
- 2020(Publication Date)
- De Gruyter(Publisher)
8 Electrochemistry in energy conversion and storage 8.1 Batteries 8.1.1 Electrochemical cell – fundamentals The simplest electrochemical cell contains two electronic conductors (electrodes, mostly metals) and ionic liquid conductor containing ionic species (electrolyte) be-tween them. The electronic conductor and its interface with electrolyte serve as the place where the electrochemical reactions occur. To prevent any unwanted reaction with electrolyte species, it is often necessary to apply diaphragm dividing the cell into two parts (half cells). In this case, the additional potential is formed, called the liquid junction potential , which can be limited by means of a salt bridge or binary electrode. The schemes of exemplary cells without (A) and with junction (B) are as follows: Ag Ag 2 O , KOH aq , HgO Hg , Zn ZnSO 4 , aq CuSO 4 , aq Cu ð A Þ Zn ZnSO 4 , aq a 1 ð Þ CuSO 4 , aq a 2 ð Þ Cu , Zn ZnCl 2 , aq a 1 ð Þ AgNO 3 , aq a 1 ð Þ Ag B ð Þ The other kind of electrochemical cells are the concentration cells, in which the two half cells differ only in the concentration of the same electroactive species taking part in electrochemical reaction: Ag j AgCl , HCl aq a 1 ð Þ , H 2 ð p 1 Þ Pt − Pt j j H 2 p 2 ð Þ , HCl aq a 2 ð Þ , AgCl j Ag A ð Þ Ag AgCl , HCl aq a 1 ð Þ HCl aq a 2 ð Þ , AgCl j Ag B ð Þ As the voltage of concentration cells is low, they are not used practically in energy production and will not be considered further. Note that a slash | represents a phase boundary, a double slash || represents the phase boundary whose potential is negligible as a result of application of separator between the two electrolytes, and coma separates two components in the same phase. The overall chemical reaction in an electrochemical cell is formed by two inde-pendent half-reactions taking part in half cells, each of them characterized by the interfacial potential difference Δ φ (Galvani potential, see Chapter 1).
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
19.7 Electrolytic Cells 971 FIGURE 19.16 Diagram of the internal parts of a photovoltaic cell. a. Glass protective covering, b. antireflective coating, c. electrode mesh, d. n-type semiconductor material, e. p-type semiconductor material, f. electrode. The thickness of this cell is approximately 0.3 mm, and the n-type semiconductor is approximately 0.002 mm thick. FIGURE 19.17 Acres of photovoltaic cells can produce large amounts of energy. Stocktrek Images/Getty Images a b c d e Sunlight f 19.7 Electrolytic Cells In our preceding discussions, we’ve examined how spontaneous redox reactions can be used to generate electrical energy. We now turn our attention to the opposite process, the use of electrical energy to force nonspontaneous redox reactions to occur. When electricity is passed through a molten (melted) ionic compound or through a solu- tion of an electrolyte, a chemical reaction occurs that we call electrolysis. A typical electroly- sis apparatus, called an electrolysis cell or electrolytic cell, is shown in Figure 19.18. This particular cell contains molten sodium chloride. (A substance undergoing electrolysis must be molten or in solution so its ions can move freely and conduction can occur.) Inertelectrodes— electrodes that won’t react with the molten NaCl or the electrolysis products—are dipped into the cell and then connected to a source of direct current (DC) electricity. The DC source serves as an “electron pump,” pulling electrons away from one electrode and pushing them through the external wiring onto the other electrode. The electrode from which electrons are removed becomes positively charged, while the other electrode becomes negatively charged. When electricity starts to flow, chemical changes begin to happen. At the positive electrode, oxidation occurs as electrons are pulled away from negatively charged chloride ions. Because of the nature of the chemical change, therefore, thepositiveelectrode becomestheanode.
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- Amos Turk(Author)
- 2013(Publication Date)
- Academic Press(Publisher)
16 & GALVANIC CELLS AND THE DRIVING FORCE OF CHEMICAL REACTIONS 16.1 • INTRODUCTION We saw in the last chapter that an electric current can cause chemical reactions. The process can be reversed, and electricity obtained from a chemical reaction. This discovery, by Alessandro Volta in 1796, pro-vided for the first time a source of continuous electric current, and thus made possible the great electrical discoveries of the early nineteenth century, including the electrochemical observations that we have de-scribed. Today, the electrochemical cell, or battery of cells, has been generally superseded by the generator as a source of electricity. How-ever, cells still have a place as portable sources of relatively small amounts of power. What is more important, from the point of view of chemistry, is that the electrical work that can be obtained from a chem-ical reaction provides a direct measure of the driving force of the re-action, and, not quite so directly, of its equilibrium constant. One of the most familiar chemical reactions is the reaction between zinc and a soluble copper salt, Zn (c) + CuS0 4 (a Cu (c) + ZnS0 4 (aq) or, in ionic form, Zn (c) + C u 2+ > Cu (c) + Z n 2+ The products in this reaction have lower enthalpy than the reactants; that is, AH is negative, or the reaction is exothermic. Merely adding a piece of zinc to a CuS0 4 solution results in the liberation of heat, but 282 2 8 3 • 16.2 E L E C T R I C I T Y F R O M A C H E M I C A L R E A C T I O N yields no work, in mechanical or electrical form, aside from a minute amount that the reaction mixture will do if it expands against atmo-spheric pressure. Of course, the liberated heat can be partially con-verted to work (as in a steam engine), but this process is notoriously inefficient and does not help us to measure the maximum available work—the quantity in which we are especially interested.
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