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Week 2 - Electrochemistry.pdf;JSESSIONID=KyvgrKcbSwGKFy1wzrddG2pCTmJyX5jNPfL9LqYnJ3B1zwvPL8mZ!-27772

Week 2 - Electrochemistry.pdf;JSESSIONID=KyvgrKcbSwGKFy1wzrddG2pCTmJyX5jNPfL9LqYnJ3B1zwvPL8mZ!-27772

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1 INDC 2003 – Instrumental Analysis Electroanalytical Chemistry: Direct Potentiometry and Ion Selective Electrodes Maria Skyllas-Kazacos School of Chemical Sciences and Engineering Electrochemical Cells Oxidation (loss of electrons) at the anode Reduction (gain of electrons) at the cathode Galvanic cells involve spontaneous reactions (negative G reaction ) Electrolytic cell require external source of energy (positive G reaction ) Battery charging is electrolytic process - requires external source of energy Battery discharging is a galvanic process - releases energy
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2 Cell Voltage G o = - nFE o cell • E o cell = E o cathode - E o anode (where E o is standard reduction potential) • ie E o cell > 0 for galvanic cell E o cell < 0 for electrolytic cell Electrode Potentials Electrochemical cells made up of 2 half-cell reactions Each half-cell reaction has a particular chemical potential or driving force Cannot measure half-cell potentials, only differences - therefore measure all half-cell potentials relative to arbitrary reference ---> standard hydrogen electrode (SHE) - ie 2 H + + 2 e -------> H 2(g) where E SHE = 0 V at all T’s I.U.P.A.C. convention: - all half-reactions written as reductions (eg Al 3+ + 3e -> Al) - Resultant half-cell potentials vs SHE will have: (a) positive sign if acts as oxidiser relative to SHE - ie reaction goes as written vs SHE (b) negative sign if acts as reducer relative to SHE
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