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The Aqueous Chemistry of Oxides$
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Bruce C. Bunker and William H. Casey

Print publication date: 2016

Print ISBN-13: 9780199384259

Published to Oxford Scholarship Online: November 2020

DOI: 10.1093/oso/9780199384259.001.0001

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PRINTED FROM OXFORD SCHOLARSHIP ONLINE (oxford.universitypressscholarship.com). (c) Copyright Oxford University Press, 2021. All Rights Reserved. An individual user may print out a PDF of a single chapter of a monograph in OSO for personal use. date: 04 December 2021

The Electrochemistry of Oxides

The Electrochemistry of Oxides

Chapter:
(p.306) 11 The Electrochemistry of Oxides
Source:
The Aqueous Chemistry of Oxides
Author(s):

Bruce C. Bunker

William H. Casey

Publisher:
Oxford University Press
DOI:10.1093/oso/9780199384259.003.0018

Many of the critical reactions considered in this book involve the addition or subtraction of protons from oxides. In this chapter, we consider another species that can change oxide charge distributions and reactivity dramatically: the electron. Many oxides contain cations that have access to more than one oxidation state in water. Cations in these oxides can either donate or accept electrons to change their charge, or oxidation state. Oxidation reactions involve the loss of electrons as they are donated to other species, resulting in an increase in the cation charge or valence, whereas reduction reactions involve the capture of an electron resulting in a decrease in the cation valence. Below, the basics of electrochemistry are first described in the context of the redox chemistry of water and representative oxide systems. Second, we describe the fundamentals of electron-transfer reactions in oxides and the impact of electron transfer on the acid–base, ion-exchange, and ligand-exchange reactions of the host oxide. Finally, we discuss the behavior of oxides in electrochemical energy storage devices and the role that nanotechnology has in optimizing electrochemical performance. Electrochemical reactions are typically written in the context of the geometry of a battery or a galvanic cell. For the cell shown in Figure 11.2, which contains metal electrodes immersed in aqueous solutions containing solvated cations, the net reaction can be written as … Cu°(s)+2Fe3+ (aq) ?2Fe2+ (aq)+Cu2+ (aq) (11.1)… In this reaction, Cu metal is being oxidized to form Cu(II), whereas Fe(III) is being reduced to form Fe(II). The individual oxidation and reduction reactions leading to Eq. 11.1 are referred to as half-reactions: Cu°(s) ?Cu2+ (aq)+2e-Eo =-0.34 V (11.2) 2Fe3+ (aq) ?2Fe2+ (aq)+ Fe 2- (aq) Eo =+0.77 V (11.3) … The net result of Eqs. 11.2 and 11.3 in the context of Figure 11.2 is the transfer of electrons from the Cu-containing compartment in which oxidation occurs, which is called the anode, to the Fe-containing compartment in which reduction occurs, which is called the cathode. This electron transfer generates an electric current, a charge-compensating ion current (in the salt bridge), as well as a voltage.

Keywords:   Faraday constant, acid rain, batteries, cathodes, fuel cells, hydrogen electrode, lead acid battery, mixed-valence oxides, oxidation potential, peroxides

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