Exam 2 Review B.notes

Exam 2 Review B.notes - Chapter 6 Electrochemistry...

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Chapter 6 Electrochemistry
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Electrochemistry: The Basics Redox Reactions: simply the coupling of reduction and oxidation half reactions. Oxidation: the element gives up electrons and its oxidation number increases. Reduction: the element accepts electrons and its oxidation number decreases. ( ) Ag e Ag s + - + 2 ( ) 2 Cu s Cu e + - + 2 2 ( ) ( ) ( ) 2 ( ) Ag aq Cu s Cu aq Ag s + + + +
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Electrochemistry: Balancing Reactions Balance the following redox reaction in an acidic solution: - - - 3 4 IO (aq) + Re(s) ReO (aq) + I (aq) )
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Electrochemistry: Driving Forces (E cell and Free Energy) o o o cell E = E - E cathode anode E red is a measure of how easily something is reduced relative to the standard hydrogen electrode. In a spontaneous reaction ( G<0), the species with the higher E red will be reduced and the one with the lower E red will be oxidized, such that E cell is always positive: Both are E red values, do not change the sign from the E red value given in the table. o = cell G nFE - o G is related to E cell by a constant:
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Electrochemistry: Non-Standard Conditions o RT cell cell nF E = E - ln(Q) The Nernst Equation is used for calculating E cell under non- standard conditions: As Q approaches K eq the E cell will approach zero; as the system approaches equilibrium, the driving force of the cell is depleted. = cell G nFE - We can also generalize to non-standard conditions:
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Electrochemistry: Setting up an Electrochemical Cell Electrolytic Cell Galvanic/Voltaic Cell The difference between these cells is that a battery is used to force the electrons to move in a
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This note was uploaded on 09/01/2008 for the course LIFE SCIEN 1B taught by Professor Wakeley during the Spring '08 term at Harvard.

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Exam 2 Review B.notes - Chapter 6 Electrochemistry...

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