zumdahl_chemprin_6e_csm_ch11

zumdahl_chemprin_6e_csm_ch11 - CHAPTER 11 ELECTROCHEMISTRY...

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413 CHAPTER 11 ELECTROCHEMISTRY Galvanic Cells, Cell Potentials, and Standard Reduction Potentials 15. Electrochemistry is the study of the interchange of chemical and electrical energy. A redox (oxidation-reduction) reaction is a reaction in which one or more electrons are transferred. In a galvanic cell, a spontaneous redox reaction occurs that produces an electric current. In an electrolytic cell, electricity is used to force a nonspontaneous redox reaction to occur. 16. Magnesium is an alkaline earth metal; Mg will oxidize to Mg 2+ . The oxidation state of hydrogen in HCl is +1. To be reduced, the oxidation state of H must decrease. The obvious choice for the hydrogen product is H 2 (g), where hydrogen has a zero oxidation state. The balanced reaction is Mg(s) + 2 HCl(aq) MgCl 2 (aq) + H 2 (g). Mg goes from the 0 to the +2 oxidation state by losing two electrons. Each H atom goes from the +1 to the 0 oxidation state by gaining one electron. Since there are two H atoms in the balanced equation, a total of two electrons are gained by the H atoms. Hence two electrons are transferred in the balanced reaction. When the electrons are transferred directly from Mg to H + , no work is obtained. In order to harness this reaction to do useful work, we must control the flow of electrons through a wire. This is accomplished by making a galvanic cell that separates the reduction reaction from the oxidation reaction in order to control the flow of electrons through a wire to produce a voltage. 17. A typical galvanic cell diagram is: Cathode (reduction) Anode (oxidation) e Cations Anions e Salt bridge
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414 CHAPTER 11 ELECTROCHEMISTRY The diagram for all cells will look like this. The contents of each half-cell will be identified for each reaction, with all concentrations at 1.0 M and partial pressures at 1.0 atm. Note that cations always flow into the cathode compartment and anions always flow into the anode compartment. This is required to keep each compartment electrically neutral. a. Reference Table 11.1 for standard reduction potentials. Remember that o cell E = E°(cathode) E°(anode); in the Solutions Guide , we will represent E°(cathode) as o c E and represent E°(anode) as o a E . Also remember that standard potentials are not multi- plied by the integer used to obtain the overall balanced equation. ( C l 2 + 2 e 2 Cl ) × 3 o c E = 1.36 V 7 H 2 O + 2 Cr 3+ Cr 2 O 7 2 + 14 H + + 6 e o a E= 1.33 V ________________________________________________________________________ 7 H 2 O(l) + 2 Cr 3+ (aq) + 3 Cl 2 (g) Cr 2 O 7 2 (aq) + 6 Cl (aq) + 14 H + (aq) o cell E = 0.03 V The contents of each compartment are: Cathode: Pt electrode; Cl 2 bubbled into solution, Cl in solution Anode: Pt electrode; Cr 3+ , H + , and Cr 2 O 7 2 in solution We need a nonreactive metal to use as the electrode in each case since all the reactants and products are in solution. Pt is the most common choice. Another possibility is graphite.
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This note was uploaded on 03/29/2011 for the course CHEM 232 taught by Professor Malambri,w during the Spring '11 term at Kentucky.

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zumdahl_chemprin_6e_csm_ch11 - CHAPTER 11 ELECTROCHEMISTRY...

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