[Lower_S._K.]_Electrochemistry._Chemical_reactions(BookZZ.org).pdf | Electrochemistry | Redox

Please download to get full document.

View again

of 40
All materials on our website are shared by users. If you have any questions about copyright issues, please report us to resolve them. We are always happy to assist you.
Information Report
Category:

Documents

Published:

Views: 3 | Pages: 40

Extension: PDF | Download: 0

Share
Related documents
Description
Electrochemistry Chemical reactions at an electrode, galvanic and electrolytic cells A Chem1 Reference Text Stephen K. Lower ã Simon Fraser University1 Table of contents 1: Chemistry and electricity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Transcript
  Electrochemistry Chemical reactions at an electrode, galvanic and electrolytic cells A Chem1 Reference Text  Stephen K. Lower ã Simon Fraser University 1 Table of contents 1:Chemistry and electricity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 Electroneutrality . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3Potential differences at interfaces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5 2:Electrochemical cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6 Transport of charge within the cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7Cell description conventions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8Electrodes and electrode reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8Standard half-cell potentials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9Reference electrodes 11 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3:Prediction and significance of cell potentials . . . . . . . . . . . . . . . . . . . . . . . . . 12 Cell potentials and the electromotive series . . . . . . . . . . . . . . . . . . . . . . . . . . 13Cell potentials and free energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13 4:The Nernst equation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 Concentration cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23Thermodynamics of galvanic cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23Analytical applications of the Nernst equation . . . . . . . . . . . . . . . . . . . . . . . 23Membrane potentials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 27 5:Batteries and fuel cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28 Primary and secondary batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28The fuel cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 30 6:Electrochemical Corrosion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 327:Electrolytic cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35 Electrolysis in aqueous solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 36Faraday's laws of electrolysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 36Industrial electrolytic processes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37 1. To contact the author, please use the Web form at http://www.chem1.com/VT_mail.html  This document covers Electrochemistry at a level appropriate for first-year college chemistry.It was last modified on 9 July 2004 .It can be downloaded from http://www.chem1.com/acad/pdf/elchem.pdf A Web-based version will be available later. 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, see http://www.chem1.com/acad/webtext/virtualtextbook.html  Page 3 1 ã Chemistry and electricity The connection between chemistry and electricity is a very old one, going back to A LESSAN-DRO  V OLTA ' S  discovery, in 1793, that electricity could be produced by placing two dissimilar metals on opposite sides of a moistened paper. In 1800, Nicholson and Carlisle, using Volta's primitive battery as a source, showed that an electric current could decompose water into oxygen and hydrogen. This was surely one of the most significant experiments in the history of 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 between them. By 1812, the Swedish chemist B ERZELIUS  could propose that all atoms are electrified, hydrogen and the metals being positive, the nonmetals negative. In electrolysis, the applied voltage was thought to overpower the attraction between these opposite charges, pulling the electrified 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 theory of G.N. L EWIS  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 to play a central role in the development of chemistry. H UMPHREY   D  AVEY   prepared the first ele-mental sodium by electrolysis of a sodium hydroxide melt. It was left to Davey's former assistant, M ICHAEL  F  ARADAY  , to show that there is a direct relation between the amount of electric charge passed through the solution and the quantity of electrolysis products. J  AMES  C LERK   M  AXWELL  immediately saw this as evidence for the “molecule of electricity”, but the world would not be receptive to the concept of the electron until the end of the century. 1.1 Electroneutrality Nature seems to strongly discourage any process that would lead to an excess of positive or negative charge in matter. Suppose, for example, that we immerse a piece of zinc metal in pure water. A small number of zinc atoms go into solution as Zn ions, leaving their electrons behind in the metal: Zn(s) →  Zn 2+  + 2 e–   As this process goes on, the electrons which remain in the zinc cause a negative charge to build up within the metal which makes it increasingly difficult for additional positive ions to leave the metallic phase. A similar buildup of positive charge in the liquid phase adds to this inhibition. Very soon, therefore, the process comes to a halt, resulting in a solution in which the concentration of Zn 2+  is still too low (around 10  –10   M  ) to be detected by ordinary chemi-cal means.  Fig. 1: Oxidation of metallic zinc in contact with water  Page 4Electrochemistry There would be no build-up of charge if the electrons could be removed from the metal as the positive ions go into solution. One way to arrange this is to drain off the excess elec-trons 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 of no more than a few volts, and corresponds to unbalances in the concentrations of oppositely charged particles that are not chemically significant. There is nothing mysterious about this prohibition, known as the electroneutrality principle  ; it is a simple consequence of the thermodynamic work required to separate opposite charges, or to bring like charges into closer 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 other process that restores electroneutrality to the two phases. A simple way to accomplish this would be to immerse the zinc in a solution of copper sulfate instead of pure water. As you will recall if you have seen this commonly-performed experiment carried out, the zinc metal quickly becomes covered with a black coating of finely-divided metallic copper. The reaction is a simple oxidation-reduction process, a transfer of two electrons from the zinc to the cop-per:Zn (s)   →  Zn 2+  + 2 e–  Cu 2+  + 2 e–  →  Cu (s) The dissolution of the zinc is no longer inhibited by a buildup of negative charge in the metal, because the excess electrons are removed from the zinc by copper ions that come into contact with it. At the same time, the solution remains electrically neutral, since for each Zn ion introduced to the solution, one Cu ion is removed. The net reaction Zn (s)  + Cu 2+   →  Zn 2+  + Cu (s) quickly goes to completion.
Recommended
View more...
We Need Your Support
Thank you for visiting our website and your interest in our free products and services. We are nonprofit website to share and download documents. To the running of this website, we need your help to support us.

Thanks to everyone for your continued support.

No, Thanks