unit 10 - chpt 18 - electrochemistry
DESCRIPTION
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. - PowerPoint PPT PresentationTRANSCRIPT
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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
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Workbook Lesson pkt
• Lesson 13 - Formal Oxidation number assignments - with examples and homework
• Lesson 28 - Balancing Redox Reactions with examples and homework
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Sect 18.2 - 18.5 (slides provided)
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Galvanic Cell schematic
Oxidation occurs at anode (vowels)
Reduction occurs at cathode (consonants)
Oxidation produces electrons, so current flows from anode to cathode.
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Types of cells
Standard Hydrogen cell platinum electrodemetal electrodes
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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
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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
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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
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K example
• Example 18.10 pg 841
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Battery - Dry Cell
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