Section17_Electrochemistry_1

Section17_Electrochemistry_1 - Electrochemistry I Galvanic...

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Page 1 of 16 Electrochemistry I Galvanic Cells, The Nernst Equation and Batteries By the end of this lecture, you will be able to: (1) Draw a galvanic cell, and account for oxidation, reduction, cathode and anode (2) Use the standard hydrogen electrode to develop a table of reduction potentials (3) Calculate ε ° cell for any redox pair (4) Relate ε ° cell to Δ G° and w max (5) Understand how ε cell depends on concentration. (6) State and apply the Nernst equation. (7) State the anodic and cathodic reactions of lead storage and dry cell batteries. (8) Draw a schematic of an H 2 /O 2 fuel cell. (9) Debate the advantages and disadvantages of the H 2 /O 2 fuel cell.
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Page 2 of 16 Galvanic (or Voltaic) Cells During redox reactions in solution, electrons are transferred directly from the reducing agent to the oxidizing agent i.e. , no useful work is done. For example: 8 H + + MnO 4 + 5 Fe 2+ Mn 2+ + 5 Fe 3+ + 4 H 2 O Reduction: 8 H + + MnO 4 + 5e Mn 2+ + 4 H 2 O Oxidation: 5(Fe 2+ Fe 3+ + e ) To permit work to be done, we need to separate the reagents and force electrons to travel through a wire. We also need to allow charge balancing, i.e. , movement of ions. This is accomplished using a salt bridge or porous plug between the compartments. Electrode at which reduction occurs is the cathode . It is the positive terminal. Here the oxidising agent is reduced.
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Page 3 of 16 Electrode at which oxidation occurs is the anode . It is the negative terminal. Here the reducing agent is oxidised. Electrons flow from the anode to the cathode until all the oxidising agent has been reduced and/or all the reducing agent has been oxidised. The force driving the electrons is known as the electromotive force (emf) of the cell, or the cell potential ( ε cell ). ε cell is measured in volts (V). 1 V = 1 J C -1 The coulomb (C) is a unit of charge . The charge on one electron is 1.6 × 10 -19 coulombs. Standard Reduction Potentials We need to be able to predict the emf for a particular cell. One approach: (i) Break down overall reaction into two half- reactions (ii) Sum the emfs for each half reaction
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Page 4 of 16 But: We cannot measure absolute emfs for half-reactions. So: We need a standard. Chemists use the
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This note was uploaded on 06/20/2008 for the course CHEM 024b taught by Professor Jones during the Winter '07 term at UWO.

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Section17_Electrochemistry_1 - Electrochemistry I Galvanic...

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