C1x elchem

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  • 1. Electrochemistry Chemical reactions at an electrode, galvanic and electrolytic cellsA Chem1 Reference TextStephen K. Lower Simon Fraser University1 Table of contents1: Chemistry and electricity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3Electroneutrality . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3Potential differences at interfaces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 52: Electrochemical cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6Transport of charge within the cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7Cell description conventions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8Electrodes and electrode reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8Standard half-cell potentials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9Reference electrodes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 113: Prediction and significance of cell potentials . . . . . . . . . . . . . . . . . . . . . . . . . 12Cell potentials and the electromotive series . . . . . . . . . . . . . . . . . . . . . . . . . . 13Cell potentials and free energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 134: The Nernst equation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18Concentration cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23Thermodynamics of galvanic cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23Analytical applications of the Nernst equation . . . . . . . . . . . . . . . . . . . . . . . 23Membrane potentials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 275: Batteries and fuel cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28The ideal and the real . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28Primary and secondary batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 29Timeline of battery development . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 30The fuel cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 316: Electrochemical Corrosion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 327: Electrolytic cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35Electrolysis in aqueous solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 36Faradays laws of electrolysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37Industrial electrolytic processes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 38 1. To contact the author, please use the Web form at http://www.chem1.com/VT_mail.html

2. The Chem1 Virtual Textbook is a collection of reference textbook chapters and tutorial units providing in-depth coverage of topics in college-level General Chemistry.For a more information on the Virtual Textbook contents, seehttp://www.chem1.com/acad/webtext/virtualtextbook.htmlThis document covers Electrochemistry at a level appropriate for rst-year college chemistry.It was last modied on 23 September 2004 .It can be downloaded from http://www.chem1.com/acad/pdf/elchem.pdfA Web-based version is available at http://www.chem1.com/acad/webtext/elchem/. 3. 1 Chemistry and electricityThe connection between chemistry and electricity is a very old one, going back to ALESSAN-DROVOLTAS discovery, in 1793, that electricity could be produced by placing two dissimilarmetals on opposite sides of a moistened paper. In 1800, Nicholson and Carlisle, using Voltasprimitive battery as a source, showed that an electric current could decompose water intooxygen and hydrogen. This was surely one of the most significant experiments in the historyof chemistry, for it implied that the atoms of hydrogen and oxygen were associated with pos-itive and negative electric charges, which must be the source of the bonding forces betweenthem. By 1812, the Swedish chemist BERZELIUS could propose that all atoms are electrified,hydrogen and the metals being positive, the nonmetals negative. In electrolysis, the appliedvoltage was thought to overpower the attraction between these opposite charges, pulling theelectrified atoms apart in the form of ions (named by Berzelius from the Greek for travel-ers). It would be almost exactly a hundred years later before the shared electron pair theoryof G.N. LEWIS could offer a significant improvement over this view of chemical bonding.Meanwhile the use of electricity as a means of bringing about chemical change continued toplay a central role in the development of chemistry. HUMPHREY DAVEY prepared the first ele-mental sodium by electrolysis of a sodium hydroxide melt. It was left to Daveys formerassistant, MICHAEL FARADAY, to show that there is a direct relation between the amount ofelectric charge passed through the solution and the quantity of electrolysis products. JAMESCLERK MAXWELL immediately saw this as evidence for the molecule of electricity, but theworld would not be receptive to the concept of the electron until the end of the century.1.1 ElectroneutralityNature seems to strongly discourage any process that would lead to an excess of positive ornegative charge in matter. Suppose, for example, that we immerse a piece of zinc metal inpure water. A small number of zinc atoms go into solution as Zn ions, leaving their electronsbehind in the metal: Zn(s) Zn2+ + 2eAs this process goes on, the electrons which remain in the zinc cause a negative charge tobuild up within the metal which makes it increasingly difficult for additional positive ions toleave the metallic phase. A similar buildup of positive charge in the liquid phase adds to thisinhibition. Very soon, therefore, the process comes to a halt, resulting in a solution in whichthe concentration of Zn2+ is still too low (around 1010 M) to be detected by ordinary chemi-cal means.Fig. 1: Oxidation of metallic zinc in contact with water Page 3 4. There would be no build-up of opposing charges in the two phases if the excess electrons could be removed from the metal or the positive ions consumed as the reaction proceeds. For example, we could drain off the electrons left behind in the zinc through an external circuit that forms part of a complete electrochemical cell; this we will describe later. Another way to remove electrons is to bring a good electron acceptor (that is, an oxidizing agent) into contact with the electrode. A suitable electron acceptor would be hydrogen ions; this is why acids attack many metals. For the very active metals such as sodium, water itself is a sufficiently good electron acceptor.The degree of charge unbalance that is allowed produces differences in electric potential ofno more than a few volts, and corresponds to unbalances in the concentrations of oppositelycharged particles that are not chemically significant. There is nothing mysterious about thisprohibition, known as the electroneutrality principle; it is a simple consequence of thethermodynamic work required to separate opposite charges, or to bring like charges intocloser contact. The additional work raises the free energy of the process, making it less spon-taneous.The only way we can get the oxidation of the metal to continue is to couple it with some otherprocess that restores electroneutrality to the two phases. A simple way to accomplish thiswould be to immerse the zinc in a solution of copper sulfate instead of pure water. As you willrecall if you have seen this commonly-performed experiment carried out, the zinc metalquickly becomes covered with a black coating of finely-divided metallic copper. The reactionis a simple oxidation-reduction process, a transfer of two electrons from the zinc to the cop-per: Zn(s) Zn2+ + 2eCu2+ + 2e Cu(s)The dissolution of the zinc is no longer inhibited by a buildup of negative charge in themetal, because the excess electrons are removed from the zinc by copper ions that come intocontact with it. At the same time, the solution remains electrically neutral, since for each Znion introduced to the solution, one Cu ion is removed. The net reactionZn(s) + Cu2+ Zn2+ + Cu(s)quickly goes to completion.Page 4 Electrochemistry 5. 1.2 Potential differences at interfacesThe transition region between two phases consists of a regionof charge unbalance known as the electric double layer. As itsname implies, this consists of an inner monomolecular layerof adsorbed water molecules and ions, and an outer diffuseregion that compensates for any local charge unbalance thatgradually merges into the completely random arrangement ofthe bulk solution. In the case of a metal immersed in purewater, the electron fluid within the metal causes the polarwater molecules to adsorb to the surface and orient them-selves so as to create two thin planes of positive and negativecharge. If the water contains dissolved ions, some of thelarger (and more polarizable) anions will loosely bond (chemi-sorb) to the metal, creating a negative inner layer which iscompensated by an excess of cations in the outer layer.Electrochemistry is the study of reactions in which chargedFig. 2: Electric double layerparticles (ions or electrons) cross the interface between twoat an electrode surfacephases of matter, typically a metallic phase (the electrode)and a conductive solution, or electrolyte. A process of this kindcan always be represented as a chemical reaction and is known generally as an electrodeprocess.Electrode processes take place within t