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Unit 10 - Chpt 18 - Electrochemistry

Unit 10 - Chpt 18 - Electrochemistry. Balance Redox equations HW set1: Chpt 18 - pg. 862-865 # 30, 32 - Due Tues. Apr 20 HW set2: Chpt 18 - pg. 862-865 # 40, 44, 50, 54, 60, 65, 74 - Due Fri. Apr 23. Workbook Lesson pkt.

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Unit 10 - Chpt 18 - Electrochemistry

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  1. Unit 10 - Chpt 18 - Electrochemistry • Balance Redox equations • HW set1: Chpt 18 - pg. 862-865 # 30, 32 - Due Tues. Apr 20 • HW set2: Chpt 18 - pg. 862-865 # 40, 44, 50, 54, 60, 65, 74 - Due Fri. Apr 23

  2. Workbook Lesson pkt • Lesson 13 - Formal Oxidation number assignments - with examples and homework • Lesson 28 - Balancing Redox Reactions with examples and homework

  3. Sect 18.2 - 18.5 (slides provided)

  4. Galvanic Cell schematic Oxidation occurs at anode (vowels) Reduction occurs at cathode (consonants) Oxidation produces electrons, so current flows from anode to cathode.

  5. Types of cells Standard Hydrogen cell platinum electrode metal electrodes

  6. Cell Potential & Nernst Equation • Galvanic Cell Potentials - free energy Go = -nFEo F is Faraday constant 96485 C/mol e- n = moles of e- from balanced Redox eqn • Concentration Cell Potentials G = Go + RTlnQ or K if Ecell = 0

  7. Cell Potential & Nernst Equation G = Go + RT ln Q -nFE = -nFEo +RT ln Q E = Eo - RT/nF ln Q E = Eo - 0.0592/n log Q So E of a cell with concentrations not equal to 1 M is the std cell potential with the correction remember to know electrons transferred in Redox

  8. Equilibrium, K constant E = Eo - RT/nF log Q At equilibrium Ecell = 0 and Q = K 0 = Eo - 0.0592/n log K log K = nEo / 0.0592

  9. K example • Example 18.10 pg 841

  10. Battery - Dry Cell

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